CHAPTER 2 Atoms, Molecules, Ions
CHAPTER 2 Atoms, Molecules, Ions
Basic postulates of Dalton's atomic theory:
By 1800, two important laws were established.
Law of conservation of mass (Lavoisier) CH4 + 2 O2
CO2 + 2 H2O
There is no change in mass when a chemical reaction occurs.
Law of definite proportions (law of constant composition): (Proust) A compound always contains elements combined in the same proportion by mass.
Lets look at examples: 1.000 g water (H2O)
0.111 g H 0.889 g O
by y mass,, H to O ratio is always y 1:8.01
1.000 g methane (CH4)
0.250 g H 0.750 g C
by mass, H to C ratio is always 1:3
ATOMIC THEORY OF MATTER John Dalton (1803-07) Developed the atomic theory of matter – it p observation. came out of experimental
Basic idea: All matter, whether element, compound, or mixture, is composed of small, indivisible particles called atoms. Atoms are the basic building blocks of matter.
1) Matter is composed of tiny particles called atoms. 2) All atoms of an element are identical in mass and other properties. 3) Atoms of different elements differ in mass and other properties.
4) Compounds are composed of atoms of different elements combined in fixed proportions by mass. The numbers of atoms of elements in a compound is a small whole number ratio. 5) Atoms are indestructible. Atoms are neither created nor destroyed in chemical reactions; p y rearranged g to yield y new substances. simply
This theory explained the law of conservation of mass and the law of constant composition.
Sir Humphrey Davy (1778-1829): One of the most renowned scientists of the time. He did not accept Dalton’s Atomic Theory. Consequently, it was not initially widely accepted.
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DALTON & PREDICTIVE POWER
There are many examples of the law of multiple proportions.
Dalton’s atomic theory predicted the law of multiple proportions.
Law of multiple proportions: When two elements combine to form more than one compound, the masses of one element that combine with a fixed mass of the other element are in ratios of small whole numbers.
This provided powerful support of the theory but it does not prove the theory. Sir Humphrey Davy is now convinced of Dalton’s theory!!
Consider compounds composed of only C and O: Fixed amount of C CO 1.00 g C
1.33 g O
CO2
2.66 g O
1.00 g C
Ratio of oxygen (O) between the two compounds: 2.66 = 2.00 1.33
(small whole number ratio)
Two compounds of sulfur (S) and oxygen (O): A
Fixed amount of S 1.00 g S
0.998 g O
B
1.00 g S
1.497 g O
Ratio of oxygen (O) between the two compounds:
1.497 = 1.5 or 3 2 (small whole number ratio) 0.998 AÎ SO2
Atom: The smallest particle of an element that retains the chemical properties of the element. After Dalton’s atomic theory (prior to 1850) it was believed that atoms were simply indivisible balls of matter, matter indestructible, indestructible unchangeable. After 1850, evidence suggested that atoms are composed of even smaller particles called
subatomic particles.
Hence, atoms have a complex structure.
Discovery of the First Subatomic Particle Phenomenon of an electrical discharge in an evacuated glass tube, gas discharge tube (Crooks Tube). Tube) Cathode ray tube (CRT) – crude television tube.
BÎ SO3
2
Cat hode (-)
high volt age
Anode ( +)
Cathode Ray Tube
f luo resc ent screen
Cathode emits an invisible ray called a cathode ray. It causes gases in the tube to glow and yields a bright spot when it strikes a fluorescent screen (basis for TV screen).
Cathode Ray Tube
Figure 2.3
J.J. THOMSON’S CRT TUBE
CRT EXPERIMENTAL OBSERVATIONS Metal plates exposed to Cathode rays acquire a negative charge. J. J. Thomson (1897): Found that magnetic and electric fields bend Cathode rays. -negatively charged particles -same properties regardless of cathode material -suggests that Cathode rays are a basic
component of all matter Cathode rays are a beam of electrons.
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J. J. THOMSON 1856-1940
RADIOACTIVITY Röntgen (1895): Observed that cathode rays caused some materials to also emit rays.
measured the charge-to-mass ratio of electron
⎛ charge ⎞ 8 ⎜ ⎟ = −1.76 X 10 C/g ⎝ mass ⎠ minus sign indicates negative charge of electron C Î Coulomb (unit of electric charge) awarded Nobel Prize in Physics for this work
emitted rays were able to pass through matter (thick boards of wood)
not deflected by magnetic i or electric l i fi fields ld Consequently, these rays were not composed of charged particles.
These mysterious rays were coined “Xrays”.
April, 1897 - date of electron discovery
Becquerel (1896): Found that some
compounds emitted rays spontaneously without
cathode ray stimulation.
MARIE CURIE 1867-1934 Suggests term "radioactivity" describes spontaneous emission of particles and/or radiation from matter. Marie Curie was awarded two Nobel prizes (1903 in physics with Becquerel for “radioactivity”); (1911 in chemistry for discovery of radium and polonium).
Figure 2.5
Robert Millikan (1909): Determined the charge on an electron with the “Oil Drop Experiment”.
-1.60 X 10-19 C
Now, electron mass can be determined:
⎛ mass of electron ⎞ h ) = mass off electron l t ⎜ ⎟(charge charge ⎝ ⎠
(
)
1g ⎛ ⎞ - 19 − 28 g ⎜ ⎟ − 1.60 X 10 C = 9.09 X 10 ⎝ - 1.76 X 10 8 C ⎠ The electron was well characterized by 1909.
ERNEST RUTHERFORD 1871-1937 Found that there are three types of rays produced by radioactivity.
Alpha (α) rays: A beam of positively charged particles (α particles or Helium nuclei).
Beta (β) rays: A beam of negatively charged particles (β particles or electrons). Gamma (γ) rays: High energy radiation; no charge; do not consist of particles, similar to X-rays.
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Radioactive Rays
Figure 2.6
STRUCTURE OF THE ATOM (Before 1908) We know that atoms – contain negatively charged electrons – are electrically neutral – also contain positively charged matter
J. J. Thomson (Early 1900’s – before 1908)
Figure 2.7
Ernest Rutherford (1910) with Geiger and Marsden • Studied scattering of a beam of α (alpha) particles by a very thin gold foil. • Expected the dense α particles to undergo only very slight scattering. • Observed small scattering of α particles, but some α particles had large scattering angles. • Some α particles actually bounced back!!!
α-particle scattering experiment
Model of the atom: -composed of a uniform sphere of positively charged matter in which negatively charged electrons are embedded -electrons l t only l comprise i a small ll fraction f ti of f atom’s mass (ca. 1/2000 the mass of an atom)
This is referred to as the “plum pudding” (English dessert) model for the structure of the atom.
Figure 2.8
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RUTHERFORD’S MODEL OF THE ATOM (1911)
• Atom is composed of two distinct parts: – 1) A dense central core (nucleus) that carries a positive charge. – 2) Negatively charged electrons surround the h nucleus l and d are far f removed d from f the h nucleus.
• The atom is mostly empty space !! • This is called the “nuclear model” of the atom; the beginning of the “modern view of atomic structure”.
α-particle scattering experiment
Figure 2.8
Rutherford’s α-particle scattering experiment
Figure 2.8
Rutherford’s α-particle scattering experiment: These results were inconsistent with Thomson’s “plum-pudding” model of atom. How ? Should only see only small scattering angles for the α-particles. Based on this new information, Rutherford postulated a new model for the structure of the atom in 1911.
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elect r o ns
_
1 oz hummingbird
dense nucleus
+ Your body (150 lbs) (hydrogen nucleus)
Modern View of the St t Structure off th the At Atom
Hummingbird is 2 miles from your body!!!
The atom is composed of three subatomic particles:
Protons: Positively charged particles that are part of the nucleus. Neutrons: Neutral particles that are part of f the th nucleus nucleus. Electrons: Negatively charged particles that are far removed from the nucleus. There are additional subatomic particles but these are not important to chemistry.
Protons: Discovered by Rutherford in 1911. Neutrons: Discovered by James Chadwick in 1932. Neutrons and protons are part of the nucleus y similar masses. and have very The neutron was difficult to discover and characterize because it is neutral (it is not affected by magnetic and electric fields as are charged particles).
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Properties of Subatomic Particles (Table 2.1) Particle Electron Proton N t Neutron
Charge -1.60 X 10-19 C +1.60 X 10-19 C 0
Mass (g) 9.109 X 10-28 g 1.673 X 10-24 g 24 g 1 675 X 10-24 1.675
Note: Proton and electron have the same magnitude charge but opposite in sign.
elect r o ns
_
1 oz hummingbird
dense nucleus
+ Your body (150 lbs) (hydrogen nucleus)
Hummingbird is 2 miles from your body!!!
For convenience: Protons +1
Electrons
-1
FACTS ABOUT ATOMS Atoms are electrically neutral, therefore: # protons = # electrons l Protons and neutrons Î much more massive than electrons (ca. a factor of 1840). Essentially all of an atoms’ mass is found in the nucleus.
FACTS ABOUT ATOMS Atoms are very small: Mass ≈ 1 X 10-23 g Diameter (atom) ≈ 1 X 10-10 m Diameter (nucleus) ≈ 1 X 10-14 m (10,000 times smaller than the atomic radius) An atom is mostly empty space!!!
What characterizes an element? The number of protons. Henry Moseley (1913) For example:
1 Proton Î H (hydrogen) 6 Protons Î C (carbon) 15 Protons Î P (phosphorus) Atomic number (Z): Number of protons, identifies the element – chemical identity.
Mass number (A): The total number of protons and neutrons for the atom. For example: Carbon: 6 protons and 6 neutrons Mass number of 12. Atoms of a particular element can have variable numbers of neutrons.
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For example: C:
6 protons + 7 neutrons = 13
C:
6 protons + 8 neutrons = 14
Neutrons only affect the mass of the atom, atom not the chemical properties.
C - 12; C - 13; C - 14 These are called isotopes.
Isotopes: Atoms of the same element that have different mass numbers Î different numbers of neutrons.
The # of neutrons only affect the mass of the atom, not the chemical properties. Atoms of the same element that have different mass numbers have different #’s of neutrons. The nucleus of a specific isotope is called a nuclide.
Chemical Symbols: Isotopes Example: Carbon Isotopes mass number atomic number
Naturally occurring
12 6
C
13 6
C
radioactive
14 6
C
These carbon isotopes have identical chemical properties.
Important Radioactive Isotopes and Their Uses • Cobalt-60 (
60Co)
– Cancer radiation therapy
• Cesium-137 Cesium 137 (
137Cs+
ion)
– Food Sterilization (as cesium chloride)
• Iodine-129 (
129I-
ion)
– Thyroid radiation therapy (as NaI)
• Technetium-99 (
99Tc)
– Medical imaging of internal organs
ISOTOPES J. J. Thomson Î Discovered isotopes (19101915) Ne-20, Ne-22 Discovery of isotopes is in conflict with Dalton's
atomic i theory h (Postulate (P l 2). 2)
Modify theory to include isotopes: –"all atoms of a given element have the same number of protons and chemical properties".
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For any atom: Mass number = # of protons + # of neutrons and # of protons = # of electrons If we know two of the former, the third can be determined. For example, how many neutrons are in the O-18 nuclide? All oxygen atoms have 8 protons. Number of neutrons = 18 – 8 = 10
19th century – Scientists recognized that some elements had very similar chemical properties. Mendeleev (Russian) and Meyer (German) 1869 Elements were arranged by increasing atomic number (mass originally) in a horizontal row. Elements with similar chemical properties arranged in columns called Groups. Proposed an extensive tabulation of the elements based on the regular, periodic recurrence of properties; known as the periodic table
• How many protons, neutrons, and electrons are in an atom of Tungsten ( 184W) ?
– Atomic number, Z = 74, Î 74 protons – For an atom, # protons = # electrons; Î 74 electrons – Mass number, A = 184, Î 184 - 74 = 110 neutrons Figure 2.10
Time-Line for Discovery of Elements
Three categories for the elements: METALS NONMETALS METALLOIDS
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Metals, Nonmetals, Metalloids
METALS • • • • • •
located left and center of Table good conductors of heat and electricity luster (shiny) ductile (drawn into wire) malleable (pounded into sheets) All metals are solids at room temperature except mercury (Hg) which is liquid.
Periodic Table
NONMETALS • located to the right of the Table • typically have properties opposite to those for metals – poor conductors of heat and electricity – solids are brittle (not malleable or ductile) – dull surface (not shinny)
Group A elements Î representative or main group elements p B elements Î transition ((metal)) Group elements Two rows below the table Î inner transition (metal) elements Lanthanide elements (58-71) Actinide elements (90-103)
Periodic Table
METALLOIDS • located to right-center of Table • divide nonmetals and metals • properties intermediate between metals and nonmetals. • Most M t elements l t are metals t l Of the first 92 elements 67 Î metals 17 Î nonmetals 8 Î metalloids
• Group numbers – 1A, 2A, 3A......8A – 1B, 2B, 3B.......8B • Some groups have special names – – – – –
Group Group Group Group Group
1A 2A 6A 7A 8A
Î Î Î Î Î
alkali lk l metals l (not ( hydrogen) h d ) alkaline earth metals chalcogens halogens noble (rare) gases
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Modern Periodic Table
Elemental Forms of Important Nonmetallic Molecules: • H2, O2, N2, F2, Cl2 - gases @ RT (room temp.) • Br2 - liquid @ RT • I2 - solid @ RT • All are diatomic molecules (made up of two atoms) • Other nonmetals have more complex forms For example: P4, S8 - solids
Group 8A elements (noble or rare gases) exist in nature as free, single atoms. Most other elements exist as molecules or ions in nature. Atoms combine in two ways: 1) By sharing electrons to yield molecules. 2) By transferring electrons to yield ions resulting in the formation of ionic compounds.
Molecules Formed by the combination of nonmetals with nonmetals or metalloids. Have a definite shape (geometry, Chapter 10) Atoms are combined in definite proportions by mass (constant composition). All are neutral species.
Elemental oxygen: Two forms Î O2 and O3 O2 and O3 are allotropes of the element oxygen. Allotropes: Different forms of an element Carbon: Three allotropes (graphite, (graphite diamond, diamond fullerenes (1985)). Sulfur: S2, S4, S6, S8 molecules
Molecular Compounds Molecules composed of two or more different nonmetals or metalloids. –HCl, CO - diatomic molecular compounds polyatomic y molecular –CO2, NH3, H2O - p compounds –SiO2, SiH4 – polyatomic molecular compounds containing a metalloid
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CHEMICAL FORMULAS Molecular, Empirical, and Structural – Denote elements in a compound by chemical symbols and the ratios in which the atoms are combined by subscripts. Molecular formula: – A chemical formula that denotes the exact number of atoms of each element in a molecule. – Examples: H2O, NH3, CH4, C2H6, C6H6 – Subscripts denote the exact number of atoms of that element in one molecule.
Ions and Ionic Compounds • Ions: Species containing a net charge by either gaining electrons (negative charge) or by losing electrons (positive charge)
– Cation: Positively charged ion, attracted to the negative cathode ((-)) in an electrolytic cell.
– Anion: Negatively charged ion, attracted to the positive anode (+) in an electrolytic cell.
• Metals: Tend to lose electrons Î cations. • Nonmetals: Tend to gain electrons Î anions.
Monatomic Ions
Empirical formula (simplest): – Gives the smallest whole-number ratio of atoms of each element in the compound. – Compare molecular and empirical formulas: Molecular M l l C6H6 C2H4 CH4 N 2O 5
Empirical E i i l CH CH2 CH4 N2O5
• Ions with only one atom; metals and nonmetals
• Monatomic metal ions (Cations) – Examples: Na+, Ca2+, K+, Mg2+, Al3+ – How are they formed ? – By the loss of one or more e- ‘s
Mn+ + ne- (n = # of electrons)
M
Na
Ca2+ + 2e-
Ca
Structural Formulas: – Show which atoms are connected (chemically bonded) together. – Examples: H – O – H – H2O H – O – O – H – H2O2 – CO2 O – C – O
N2H4
H
H N N H
H
Na+ + e-
• Monatomic nonmetal ions (Anions) – – – –
F -, Cl -, O2-, N3How are they formed ? By gaining one or more e- s Examples:
X + ne-
Xn-
F + e-
F-
N + 3e-
N3-
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Common Stable Monatomic Ions & Charges
N Î 7 protons and 7 electrons N3- Î 7 protons and 10 electrons (7 + 3) Mg Î 12 protons and 12 electrons M 2+ Î 12 protons Mg t and d 10 electrons l t (12 – 2) S Î 16 protons and 16 electrons S2- Î 16 protons and 18 electrons (16 + 2)
The periodic table is very useful for determining the charges for stable monatomic ions.
Predict Charges for Stable Monatomic Ions The number of electrons Group 1A or 2A metal atoms loose is equal to the Group
number.
Group 1A: Lose
1e-
to form
M+
ions.
Polyatomic Ions Ions that consist of atoms chemically bonded together in a molecular sense and carry a net charge (positive or negative).
Examples: E l NO3 , SO42-,
PO43-,
CO32-,
NH4+
Group 2A: Lose 2e- to form M2+ ions.
The number of electrons that a nonmetal atom gains is equal to 8 minus the Group number. • Group 7A: Gain 1e- to form X- ions (8 – 7 = 1) • Group 6A: Gain 2e- to form X2- ions (8 – 6 = 2) • Group 5A: Gain 3e- to form X3- ions (8 – 5 = 3)
• Figure 2.11 (Page 54) – Common stable monatomic ions and their charges-KNOW!!!
Ionic Compounds • Composed of cations and anions interacting in appropriate proportions to yield an overall neutral compound. • e.g., common table salt – ionic compound composed of Na cations and Cl anions • A crystal lattice structure is formed.
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Cl-
Na+
Cl-
Na+
Na+ Cl-
Na+ Cl-
Cl-
Cl-
Na+
Na+ Cl-
Na+
Na+ Cl-
Every Na+ ion is surrounded by Cl- ions and every Cl- ion is surrounded by Na+ ions.
Opposite charges attract !!!
Ionic Compounds (NaCl) • An extended 3-d array is formed where the number of Na cations equals the number of Cl anions. • Consequently, discrete molecular units are not formed (i.e., we do not have individual NaCl molecules).
CHEMICAL FORMULAS OF IONIC COMPOUNDS • Formulas represent the smallest wholenumber ratio of the ions in the compound (Empirical formula). • All ionic compounds are represented by empirical formulas (formula unit). unit) • Always electrically neutral, therefore, the amount of positive charge must equal amount of negative charge.
+
=
Write Formulas For Ionic Compounds • 1) Ionic compounds are always neutral. • 2) Cation is always written first then anion. • 3) Subscripts denote number of ions of each type required to yield a neutral compound (smallest whole-number ratio); referred to as a formula f l unit. i MAKING AN IONIC COMPOUND
Al3+ + O2-
Na + Cl
NaCl
-
IONIC COMPOUND
FORMATION OF AN IONIC COMPOUND
Al3+ + O2-
Al2O3
For charge balance, the subscript of the cation is equal to the charge of the anion; the subscrip pt of the anion is equal q to the charge g of the cation. 2 (3+ charges) + 3 (2- charges) = zero or a neutral compound
Na+ + CO32-
Na2CO3
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Mg2+ + O2-
Mg2O2 INCORRECT
MgO M O iis the h smallest ll whole-number h l b ratio of ions - simplify!!!
NAMING INORGANIC COMPOUNDS We need systematic rules for naming compounds.
Four categories g of f compounds: p 1) 2) 3) 4)
Ionic compounds Acids Molecular compounds Hydrates
Ionic Compounds: Named from the IONS Write the Chemical Formulas for the ionic compounds formed from the following cations and anions: A.)
K+ and Br-
B.)
Ca2+ and NO3-
C.)
Ag+ and SO42-
D.)
Fe3+ and SO42-
MONATOMIC METAL CATIONS – Named from the elements – Examples:
Na+ Î sodium ion Ca2+ Î calcium ion Mg2+ Î magnesium ion
Naming Inorganic Compounds (Chemical Nomenclature)
Formula
Name
Over 13 million compounds are known. Some have common names. names
Examples:
Cation- first Anion- second
H2O Î water NH3 Î ammonia AsH3 Î arsine SiH4 Î silane
TRANSITION METAL IONS (Group B elements)
Can form stable ions with variable charges. For example, p Iron forms ions with variable charge:
Fe2+ and Fe3+ Use the Stock system to name these ions.
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Stock System: element name(charge) ion Roman numerals
Fe2+ Î iron(II) ion Fe3+ Î iron(III) ion
Polyatomic Anions: Many contain oxygen; classified as oxoanions; have suffixes -ate and -ite. Formula OHCNO22CO32HCO3MnO4CrO42Cr2O72PO43-
Ion Name hydroxide ion cyanide ion peroxide ion carbonate ion hydrogen carbonate ion (bicarbonate ion) permanganate ion chromate ion dichromate ion phosphate ion
Polyatomic Anions (cont.) POLYATOMIC CATIONS
Formula
Two common ions.
NH4+ - ammonium ion H3O+ - hydronium ion Name ends in –ium plus ion.
SCN C2H3O2NO2NO32 SO O32SO42ClO2ClO3-
Ion Name
thiocyanate ion acetate ion nitrite ion nitrate ion sulf lfite i ion i sulfate ion chlorite ion chlorate ion
Table 2.3 Polyatomic ions and names – Know!!!
ANION NAMES Monatomic nonmetal anions Know Table 2.2 Name derived from stem (root) of element name followed by the suffix –ide plus ion.
Ion F-
Element name fluorine
Ion Name fluoride ion
Cl-
chlorine
chloride ion
O2-
oxygen
oxide ion
S2-
sulfur
sulfide ion
Oxoanions: For a homologous series (e.g., NO3-, NO2-) – same central atom, variable number of oxygen atoms:
– ite suffix - anion with less oxygen – ate suffix - anion with more oxygen – If more than two oxyanions in series, use prefixes:
hypo- (even less oxygen than -ite) per- (even more oxygen than -ate)
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Naming ACIDS Acids: Compounds that yield hydronium ions, H3O+, when dissolved in water.
Example: ClOClO l 2ClO3ClO4-
hypochlorite ion chlor hl ite ion chlorate ion perchlorate ion
Cation named first, followed by anion: NaCl CaCl2 Ca(NO3)2 BaSO4 Al2(SO4)3
sodium chloride calcium chloride calcium nitrate barium sulfate aluminum sulfate
Note: Ion is not included in the name.
Transition metal ionic compounds Î specify charge by using the Stock system. FeO; Fe2+ and O2-;
iron(II) oxide
Fe2O3; Fe3+ and O2-;
iron(III) oxide
FeCl3; Fe3+ and Cl-;
iron(III) chloride
FeSO4; Fe2+
iron(II) sulfate
Fe2(SO4)3; Fe3+
iron(III) sulfate
H3O+ + NO3-
HNO3 + H2O Short-hand:
HNO3
Naming Ionic Compounds
H3O+ + Cl-
HCl + H2O
H2O
H+ + NO3-
Naming Acids: Acids based on anions whose name ends in -ide: Add hydro- prefix Change -ide suffix to -ic acid Cl- (chloride ion)
HCl (hydrochloric acid)
CN- (cyanide ion)
HCN (hydrocyanic acid)
S2- (sulfide ion)
H2S (hydrosulfuric acid)
Acids from oxoanions: Based on anion root name; simply change the suffix:
–ic acid – For -ite suffix; change to –ous acid
– For -ate suffix; change to
Examples: SO42- (sulfate ion) SO32- (sulfite ion) NO3- (nitrate ion) NO2- (nitrite ion) ClO4- (perchlorate ion)
H2SO4 (sulfuric acid) H2SO3 (sulfurous acid) HNO3 (nitric acid) HNO2 (nitrous acid) HClO4 (perchloric acid)
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Binary Molecular Compounds
Complication: Acids whose anion ends in ide.
Binary compounds: Contain only two elements
If these compounds are not in water (i.e., pure compounds), then we name them as binary molecular compounds.
Rules similar to those for binary ionic
Cl- (chloride ion)
HCl (hydrogen chloride) in water (hydrochloric acid)
CN- (cyanide ion)
HCN (hydrogen cyanide) in water (hydrocyanic acid)
S2- (sulfide ion)
H2S (hydrogen sulfide) in water (hydrosulfuric acid)
compounds, d e.g., N NaCl Cl (sodium ( di chloride). hl id ) Given Formula: Name first as the element,
the second as an anion (root with –ide suffix).
Specify number of atoms of each element by using Greek prefixes (Table 2.4): Greek Prefixes Prefix MonoDiDi TriTetraPentaHexa-
Meaning 1 2 3 4 5 6
Prefix mono- never used for the first element.
When the prefix ends in a or name begins with a vowel, the prefix is often omitted. Formula CO CO2 PCl3 NO2 N2O5 SO3 SO2
o and the anion a or o of the
Name carbon monoxide carbon dioxide d d phosphorus trichloride nitrogen dioxide dinitrogen pentoxide sulfur trioxide sulfur dioxide
Hydrates Hydrate: A compound with a specific number of water molecules that are strongly adhered. CuSO4 Î copper(II) sulfate (anhydrous) C SO4 . 5H2O Î copper(II) CuSO (II) sulfate lf t pentahydrate t h d t BaCl2 Î barium chloride (anhydrous) BaCl2 . 2H2O Î barium chloride dihydrate Anhydrous: Dry, without water
What are the chemical names or formulas of the following compounds ? (NH4)2SO4 Hydroiodic acid (hydrogen iodide) KMnO4 Mercury(II) nitrite Fe(NO3)3 Potassium dichromate
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Note: An ion and the corresponding neutral species have completely different chemical properties. NO2
NO2-
SO3
SO32-
Na
Na+
Fe
Fe3+
H
H+
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Basic postulates of Dalton's atomic theory:
By 1800, two important laws were established.
Law of conservation of mass (Lavoisier) CH4 + 2 O2
CO2 + 2 H2O
There is no change in mass when a chemical reaction occurs.
Law of definite proportions (law of constant composition): (Proust) A compound always contains elements combined in the same proportion by mass.
Lets look at examples: 1.000 g water (H2O)
0.111 g H 0.889 g O
by y mass,, H to O ratio is always y 1:8.01
1.000 g methane (CH4)
0.250 g H 0.750 g C
by mass, H to C ratio is always 1:3
ATOMIC THEORY OF MATTER John Dalton (1803-07) Developed the atomic theory of matter – it p observation. came out of experimental
Basic idea: All matter, whether element, compound, or mixture, is composed of small, indivisible particles called atoms. Atoms are the basic building blocks of matter.
1) Matter is composed of tiny particles called atoms. 2) All atoms of an element are identical in mass and other properties. 3) Atoms of different elements differ in mass and other properties.
4) Compounds are composed of atoms of different elements combined in fixed proportions by mass. The numbers of atoms of elements in a compound is a small whole number ratio. 5) Atoms are indestructible. Atoms are neither created nor destroyed in chemical reactions; p y rearranged g to yield y new substances. simply
This theory explained the law of conservation of mass and the law of constant composition.
Sir Humphrey Davy (1778-1829): One of the most renowned scientists of the time. He did not accept Dalton’s Atomic Theory. Consequently, it was not initially widely accepted.
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DALTON & PREDICTIVE POWER
There are many examples of the law of multiple proportions.
Dalton’s atomic theory predicted the law of multiple proportions.
Law of multiple proportions: When two elements combine to form more than one compound, the masses of one element that combine with a fixed mass of the other element are in ratios of small whole numbers.
This provided powerful support of the theory but it does not prove the theory. Sir Humphrey Davy is now convinced of Dalton’s theory!!
Consider compounds composed of only C and O: Fixed amount of C CO 1.00 g C
1.33 g O
CO2
2.66 g O
1.00 g C
Ratio of oxygen (O) between the two compounds: 2.66 = 2.00 1.33
(small whole number ratio)
Two compounds of sulfur (S) and oxygen (O): A
Fixed amount of S 1.00 g S
0.998 g O
B
1.00 g S
1.497 g O
Ratio of oxygen (O) between the two compounds:
1.497 = 1.5 or 3 2 (small whole number ratio) 0.998 AÎ SO2
Atom: The smallest particle of an element that retains the chemical properties of the element. After Dalton’s atomic theory (prior to 1850) it was believed that atoms were simply indivisible balls of matter, matter indestructible, indestructible unchangeable. After 1850, evidence suggested that atoms are composed of even smaller particles called
subatomic particles.
Hence, atoms have a complex structure.
Discovery of the First Subatomic Particle Phenomenon of an electrical discharge in an evacuated glass tube, gas discharge tube (Crooks Tube). Tube) Cathode ray tube (CRT) – crude television tube.
BÎ SO3
2
Cat hode (-)
high volt age
Anode ( +)
Cathode Ray Tube
f luo resc ent screen
Cathode emits an invisible ray called a cathode ray. It causes gases in the tube to glow and yields a bright spot when it strikes a fluorescent screen (basis for TV screen).
Cathode Ray Tube
Figure 2.3
J.J. THOMSON’S CRT TUBE
CRT EXPERIMENTAL OBSERVATIONS Metal plates exposed to Cathode rays acquire a negative charge. J. J. Thomson (1897): Found that magnetic and electric fields bend Cathode rays. -negatively charged particles -same properties regardless of cathode material -suggests that Cathode rays are a basic
component of all matter Cathode rays are a beam of electrons.
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J. J. THOMSON 1856-1940
RADIOACTIVITY Röntgen (1895): Observed that cathode rays caused some materials to also emit rays.
measured the charge-to-mass ratio of electron
⎛ charge ⎞ 8 ⎜ ⎟ = −1.76 X 10 C/g ⎝ mass ⎠ minus sign indicates negative charge of electron C Î Coulomb (unit of electric charge) awarded Nobel Prize in Physics for this work
emitted rays were able to pass through matter (thick boards of wood)
not deflected by magnetic i or electric l i fi fields ld Consequently, these rays were not composed of charged particles.
These mysterious rays were coined “Xrays”.
April, 1897 - date of electron discovery
Becquerel (1896): Found that some
compounds emitted rays spontaneously without
cathode ray stimulation.
MARIE CURIE 1867-1934 Suggests term "radioactivity" describes spontaneous emission of particles and/or radiation from matter. Marie Curie was awarded two Nobel prizes (1903 in physics with Becquerel for “radioactivity”); (1911 in chemistry for discovery of radium and polonium).
Figure 2.5
Robert Millikan (1909): Determined the charge on an electron with the “Oil Drop Experiment”.
-1.60 X 10-19 C
Now, electron mass can be determined:
⎛ mass of electron ⎞ h ) = mass off electron l t ⎜ ⎟(charge charge ⎝ ⎠
(
)
1g ⎛ ⎞ - 19 − 28 g ⎜ ⎟ − 1.60 X 10 C = 9.09 X 10 ⎝ - 1.76 X 10 8 C ⎠ The electron was well characterized by 1909.
ERNEST RUTHERFORD 1871-1937 Found that there are three types of rays produced by radioactivity.
Alpha (α) rays: A beam of positively charged particles (α particles or Helium nuclei).
Beta (β) rays: A beam of negatively charged particles (β particles or electrons). Gamma (γ) rays: High energy radiation; no charge; do not consist of particles, similar to X-rays.
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Radioactive Rays
Figure 2.6
STRUCTURE OF THE ATOM (Before 1908) We know that atoms – contain negatively charged electrons – are electrically neutral – also contain positively charged matter
J. J. Thomson (Early 1900’s – before 1908)
Figure 2.7
Ernest Rutherford (1910) with Geiger and Marsden • Studied scattering of a beam of α (alpha) particles by a very thin gold foil. • Expected the dense α particles to undergo only very slight scattering. • Observed small scattering of α particles, but some α particles had large scattering angles. • Some α particles actually bounced back!!!
α-particle scattering experiment
Model of the atom: -composed of a uniform sphere of positively charged matter in which negatively charged electrons are embedded -electrons l t only l comprise i a small ll fraction f ti of f atom’s mass (ca. 1/2000 the mass of an atom)
This is referred to as the “plum pudding” (English dessert) model for the structure of the atom.
Figure 2.8
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RUTHERFORD’S MODEL OF THE ATOM (1911)
• Atom is composed of two distinct parts: – 1) A dense central core (nucleus) that carries a positive charge. – 2) Negatively charged electrons surround the h nucleus l and d are far f removed d from f the h nucleus.
• The atom is mostly empty space !! • This is called the “nuclear model” of the atom; the beginning of the “modern view of atomic structure”.
α-particle scattering experiment
Figure 2.8
Rutherford’s α-particle scattering experiment
Figure 2.8
Rutherford’s α-particle scattering experiment: These results were inconsistent with Thomson’s “plum-pudding” model of atom. How ? Should only see only small scattering angles for the α-particles. Based on this new information, Rutherford postulated a new model for the structure of the atom in 1911.
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elect r o ns
_
1 oz hummingbird
dense nucleus
+ Your body (150 lbs) (hydrogen nucleus)
Modern View of the St t Structure off th the At Atom
Hummingbird is 2 miles from your body!!!
The atom is composed of three subatomic particles:
Protons: Positively charged particles that are part of the nucleus. Neutrons: Neutral particles that are part of f the th nucleus nucleus. Electrons: Negatively charged particles that are far removed from the nucleus. There are additional subatomic particles but these are not important to chemistry.
Protons: Discovered by Rutherford in 1911. Neutrons: Discovered by James Chadwick in 1932. Neutrons and protons are part of the nucleus y similar masses. and have very The neutron was difficult to discover and characterize because it is neutral (it is not affected by magnetic and electric fields as are charged particles).
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Properties of Subatomic Particles (Table 2.1) Particle Electron Proton N t Neutron
Charge -1.60 X 10-19 C +1.60 X 10-19 C 0
Mass (g) 9.109 X 10-28 g 1.673 X 10-24 g 24 g 1 675 X 10-24 1.675
Note: Proton and electron have the same magnitude charge but opposite in sign.
elect r o ns
_
1 oz hummingbird
dense nucleus
+ Your body (150 lbs) (hydrogen nucleus)
Hummingbird is 2 miles from your body!!!
For convenience: Protons +1
Electrons
-1
FACTS ABOUT ATOMS Atoms are electrically neutral, therefore: # protons = # electrons l Protons and neutrons Î much more massive than electrons (ca. a factor of 1840). Essentially all of an atoms’ mass is found in the nucleus.
FACTS ABOUT ATOMS Atoms are very small: Mass ≈ 1 X 10-23 g Diameter (atom) ≈ 1 X 10-10 m Diameter (nucleus) ≈ 1 X 10-14 m (10,000 times smaller than the atomic radius) An atom is mostly empty space!!!
What characterizes an element? The number of protons. Henry Moseley (1913) For example:
1 Proton Î H (hydrogen) 6 Protons Î C (carbon) 15 Protons Î P (phosphorus) Atomic number (Z): Number of protons, identifies the element – chemical identity.
Mass number (A): The total number of protons and neutrons for the atom. For example: Carbon: 6 protons and 6 neutrons Mass number of 12. Atoms of a particular element can have variable numbers of neutrons.
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For example: C:
6 protons + 7 neutrons = 13
C:
6 protons + 8 neutrons = 14
Neutrons only affect the mass of the atom, atom not the chemical properties.
C - 12; C - 13; C - 14 These are called isotopes.
Isotopes: Atoms of the same element that have different mass numbers Î different numbers of neutrons.
The # of neutrons only affect the mass of the atom, not the chemical properties. Atoms of the same element that have different mass numbers have different #’s of neutrons. The nucleus of a specific isotope is called a nuclide.
Chemical Symbols: Isotopes Example: Carbon Isotopes mass number atomic number
Naturally occurring
12 6
C
13 6
C
radioactive
14 6
C
These carbon isotopes have identical chemical properties.
Important Radioactive Isotopes and Their Uses • Cobalt-60 (
60Co)
– Cancer radiation therapy
• Cesium-137 Cesium 137 (
137Cs+
ion)
– Food Sterilization (as cesium chloride)
• Iodine-129 (
129I-
ion)
– Thyroid radiation therapy (as NaI)
• Technetium-99 (
99Tc)
– Medical imaging of internal organs
ISOTOPES J. J. Thomson Î Discovered isotopes (19101915) Ne-20, Ne-22 Discovery of isotopes is in conflict with Dalton's
atomic i theory h (Postulate (P l 2). 2)
Modify theory to include isotopes: –"all atoms of a given element have the same number of protons and chemical properties".
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For any atom: Mass number = # of protons + # of neutrons and # of protons = # of electrons If we know two of the former, the third can be determined. For example, how many neutrons are in the O-18 nuclide? All oxygen atoms have 8 protons. Number of neutrons = 18 – 8 = 10
19th century – Scientists recognized that some elements had very similar chemical properties. Mendeleev (Russian) and Meyer (German) 1869 Elements were arranged by increasing atomic number (mass originally) in a horizontal row. Elements with similar chemical properties arranged in columns called Groups. Proposed an extensive tabulation of the elements based on the regular, periodic recurrence of properties; known as the periodic table
• How many protons, neutrons, and electrons are in an atom of Tungsten ( 184W) ?
– Atomic number, Z = 74, Î 74 protons – For an atom, # protons = # electrons; Î 74 electrons – Mass number, A = 184, Î 184 - 74 = 110 neutrons Figure 2.10
Time-Line for Discovery of Elements
Three categories for the elements: METALS NONMETALS METALLOIDS
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Metals, Nonmetals, Metalloids
METALS • • • • • •
located left and center of Table good conductors of heat and electricity luster (shiny) ductile (drawn into wire) malleable (pounded into sheets) All metals are solids at room temperature except mercury (Hg) which is liquid.
Periodic Table
NONMETALS • located to the right of the Table • typically have properties opposite to those for metals – poor conductors of heat and electricity – solids are brittle (not malleable or ductile) – dull surface (not shinny)
Group A elements Î representative or main group elements p B elements Î transition ((metal)) Group elements Two rows below the table Î inner transition (metal) elements Lanthanide elements (58-71) Actinide elements (90-103)
Periodic Table
METALLOIDS • located to right-center of Table • divide nonmetals and metals • properties intermediate between metals and nonmetals. • Most M t elements l t are metals t l Of the first 92 elements 67 Î metals 17 Î nonmetals 8 Î metalloids
• Group numbers – 1A, 2A, 3A......8A – 1B, 2B, 3B.......8B • Some groups have special names – – – – –
Group Group Group Group Group
1A 2A 6A 7A 8A
Î Î Î Î Î
alkali lk l metals l (not ( hydrogen) h d ) alkaline earth metals chalcogens halogens noble (rare) gases
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Modern Periodic Table
Elemental Forms of Important Nonmetallic Molecules: • H2, O2, N2, F2, Cl2 - gases @ RT (room temp.) • Br2 - liquid @ RT • I2 - solid @ RT • All are diatomic molecules (made up of two atoms) • Other nonmetals have more complex forms For example: P4, S8 - solids
Group 8A elements (noble or rare gases) exist in nature as free, single atoms. Most other elements exist as molecules or ions in nature. Atoms combine in two ways: 1) By sharing electrons to yield molecules. 2) By transferring electrons to yield ions resulting in the formation of ionic compounds.
Molecules Formed by the combination of nonmetals with nonmetals or metalloids. Have a definite shape (geometry, Chapter 10) Atoms are combined in definite proportions by mass (constant composition). All are neutral species.
Elemental oxygen: Two forms Î O2 and O3 O2 and O3 are allotropes of the element oxygen. Allotropes: Different forms of an element Carbon: Three allotropes (graphite, (graphite diamond, diamond fullerenes (1985)). Sulfur: S2, S4, S6, S8 molecules
Molecular Compounds Molecules composed of two or more different nonmetals or metalloids. –HCl, CO - diatomic molecular compounds polyatomic y molecular –CO2, NH3, H2O - p compounds –SiO2, SiH4 – polyatomic molecular compounds containing a metalloid
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CHEMICAL FORMULAS Molecular, Empirical, and Structural – Denote elements in a compound by chemical symbols and the ratios in which the atoms are combined by subscripts. Molecular formula: – A chemical formula that denotes the exact number of atoms of each element in a molecule. – Examples: H2O, NH3, CH4, C2H6, C6H6 – Subscripts denote the exact number of atoms of that element in one molecule.
Ions and Ionic Compounds • Ions: Species containing a net charge by either gaining electrons (negative charge) or by losing electrons (positive charge)
– Cation: Positively charged ion, attracted to the negative cathode ((-)) in an electrolytic cell.
– Anion: Negatively charged ion, attracted to the positive anode (+) in an electrolytic cell.
• Metals: Tend to lose electrons Î cations. • Nonmetals: Tend to gain electrons Î anions.
Monatomic Ions
Empirical formula (simplest): – Gives the smallest whole-number ratio of atoms of each element in the compound. – Compare molecular and empirical formulas: Molecular M l l C6H6 C2H4 CH4 N 2O 5
Empirical E i i l CH CH2 CH4 N2O5
• Ions with only one atom; metals and nonmetals
• Monatomic metal ions (Cations) – Examples: Na+, Ca2+, K+, Mg2+, Al3+ – How are they formed ? – By the loss of one or more e- ‘s
Mn+ + ne- (n = # of electrons)
M
Na
Ca2+ + 2e-
Ca
Structural Formulas: – Show which atoms are connected (chemically bonded) together. – Examples: H – O – H – H2O H – O – O – H – H2O2 – CO2 O – C – O
N2H4
H
H N N H
H
Na+ + e-
• Monatomic nonmetal ions (Anions) – – – –
F -, Cl -, O2-, N3How are they formed ? By gaining one or more e- s Examples:
X + ne-
Xn-
F + e-
F-
N + 3e-
N3-
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Common Stable Monatomic Ions & Charges
N Î 7 protons and 7 electrons N3- Î 7 protons and 10 electrons (7 + 3) Mg Î 12 protons and 12 electrons M 2+ Î 12 protons Mg t and d 10 electrons l t (12 – 2) S Î 16 protons and 16 electrons S2- Î 16 protons and 18 electrons (16 + 2)
The periodic table is very useful for determining the charges for stable monatomic ions.
Predict Charges for Stable Monatomic Ions The number of electrons Group 1A or 2A metal atoms loose is equal to the Group
number.
Group 1A: Lose
1e-
to form
M+
ions.
Polyatomic Ions Ions that consist of atoms chemically bonded together in a molecular sense and carry a net charge (positive or negative).
Examples: E l NO3 , SO42-,
PO43-,
CO32-,
NH4+
Group 2A: Lose 2e- to form M2+ ions.
The number of electrons that a nonmetal atom gains is equal to 8 minus the Group number. • Group 7A: Gain 1e- to form X- ions (8 – 7 = 1) • Group 6A: Gain 2e- to form X2- ions (8 – 6 = 2) • Group 5A: Gain 3e- to form X3- ions (8 – 5 = 3)
• Figure 2.11 (Page 54) – Common stable monatomic ions and their charges-KNOW!!!
Ionic Compounds • Composed of cations and anions interacting in appropriate proportions to yield an overall neutral compound. • e.g., common table salt – ionic compound composed of Na cations and Cl anions • A crystal lattice structure is formed.
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Cl-
Na+
Cl-
Na+
Na+ Cl-
Na+ Cl-
Cl-
Cl-
Na+
Na+ Cl-
Na+
Na+ Cl-
Every Na+ ion is surrounded by Cl- ions and every Cl- ion is surrounded by Na+ ions.
Opposite charges attract !!!
Ionic Compounds (NaCl) • An extended 3-d array is formed where the number of Na cations equals the number of Cl anions. • Consequently, discrete molecular units are not formed (i.e., we do not have individual NaCl molecules).
CHEMICAL FORMULAS OF IONIC COMPOUNDS • Formulas represent the smallest wholenumber ratio of the ions in the compound (Empirical formula). • All ionic compounds are represented by empirical formulas (formula unit). unit) • Always electrically neutral, therefore, the amount of positive charge must equal amount of negative charge.
+
=
Write Formulas For Ionic Compounds • 1) Ionic compounds are always neutral. • 2) Cation is always written first then anion. • 3) Subscripts denote number of ions of each type required to yield a neutral compound (smallest whole-number ratio); referred to as a formula f l unit. i MAKING AN IONIC COMPOUND
Al3+ + O2-
Na + Cl
NaCl
-
IONIC COMPOUND
FORMATION OF AN IONIC COMPOUND
Al3+ + O2-
Al2O3
For charge balance, the subscript of the cation is equal to the charge of the anion; the subscrip pt of the anion is equal q to the charge g of the cation. 2 (3+ charges) + 3 (2- charges) = zero or a neutral compound
Na+ + CO32-
Na2CO3
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Mg2+ + O2-
Mg2O2 INCORRECT
MgO M O iis the h smallest ll whole-number h l b ratio of ions - simplify!!!
NAMING INORGANIC COMPOUNDS We need systematic rules for naming compounds.
Four categories g of f compounds: p 1) 2) 3) 4)
Ionic compounds Acids Molecular compounds Hydrates
Ionic Compounds: Named from the IONS Write the Chemical Formulas for the ionic compounds formed from the following cations and anions: A.)
K+ and Br-
B.)
Ca2+ and NO3-
C.)
Ag+ and SO42-
D.)
Fe3+ and SO42-
MONATOMIC METAL CATIONS – Named from the elements – Examples:
Na+ Î sodium ion Ca2+ Î calcium ion Mg2+ Î magnesium ion
Naming Inorganic Compounds (Chemical Nomenclature)
Formula
Name
Over 13 million compounds are known. Some have common names. names
Examples:
Cation- first Anion- second
H2O Î water NH3 Î ammonia AsH3 Î arsine SiH4 Î silane
TRANSITION METAL IONS (Group B elements)
Can form stable ions with variable charges. For example, p Iron forms ions with variable charge:
Fe2+ and Fe3+ Use the Stock system to name these ions.
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Stock System: element name(charge) ion Roman numerals
Fe2+ Î iron(II) ion Fe3+ Î iron(III) ion
Polyatomic Anions: Many contain oxygen; classified as oxoanions; have suffixes -ate and -ite. Formula OHCNO22CO32HCO3MnO4CrO42Cr2O72PO43-
Ion Name hydroxide ion cyanide ion peroxide ion carbonate ion hydrogen carbonate ion (bicarbonate ion) permanganate ion chromate ion dichromate ion phosphate ion
Polyatomic Anions (cont.) POLYATOMIC CATIONS
Formula
Two common ions.
NH4+ - ammonium ion H3O+ - hydronium ion Name ends in –ium plus ion.
SCN C2H3O2NO2NO32 SO O32SO42ClO2ClO3-
Ion Name
thiocyanate ion acetate ion nitrite ion nitrate ion sulf lfite i ion i sulfate ion chlorite ion chlorate ion
Table 2.3 Polyatomic ions and names – Know!!!
ANION NAMES Monatomic nonmetal anions Know Table 2.2 Name derived from stem (root) of element name followed by the suffix –ide plus ion.
Ion F-
Element name fluorine
Ion Name fluoride ion
Cl-
chlorine
chloride ion
O2-
oxygen
oxide ion
S2-
sulfur
sulfide ion
Oxoanions: For a homologous series (e.g., NO3-, NO2-) – same central atom, variable number of oxygen atoms:
– ite suffix - anion with less oxygen – ate suffix - anion with more oxygen – If more than two oxyanions in series, use prefixes:
hypo- (even less oxygen than -ite) per- (even more oxygen than -ate)
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Naming ACIDS Acids: Compounds that yield hydronium ions, H3O+, when dissolved in water.
Example: ClOClO l 2ClO3ClO4-
hypochlorite ion chlor hl ite ion chlorate ion perchlorate ion
Cation named first, followed by anion: NaCl CaCl2 Ca(NO3)2 BaSO4 Al2(SO4)3
sodium chloride calcium chloride calcium nitrate barium sulfate aluminum sulfate
Note: Ion is not included in the name.
Transition metal ionic compounds Î specify charge by using the Stock system. FeO; Fe2+ and O2-;
iron(II) oxide
Fe2O3; Fe3+ and O2-;
iron(III) oxide
FeCl3; Fe3+ and Cl-;
iron(III) chloride
FeSO4; Fe2+
iron(II) sulfate
Fe2(SO4)3; Fe3+
iron(III) sulfate
H3O+ + NO3-
HNO3 + H2O Short-hand:
HNO3
Naming Ionic Compounds
H3O+ + Cl-
HCl + H2O
H2O
H+ + NO3-
Naming Acids: Acids based on anions whose name ends in -ide: Add hydro- prefix Change -ide suffix to -ic acid Cl- (chloride ion)
HCl (hydrochloric acid)
CN- (cyanide ion)
HCN (hydrocyanic acid)
S2- (sulfide ion)
H2S (hydrosulfuric acid)
Acids from oxoanions: Based on anion root name; simply change the suffix:
–ic acid – For -ite suffix; change to –ous acid
– For -ate suffix; change to
Examples: SO42- (sulfate ion) SO32- (sulfite ion) NO3- (nitrate ion) NO2- (nitrite ion) ClO4- (perchlorate ion)
H2SO4 (sulfuric acid) H2SO3 (sulfurous acid) HNO3 (nitric acid) HNO2 (nitrous acid) HClO4 (perchloric acid)
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Binary Molecular Compounds
Complication: Acids whose anion ends in ide.
Binary compounds: Contain only two elements
If these compounds are not in water (i.e., pure compounds), then we name them as binary molecular compounds.
Rules similar to those for binary ionic
Cl- (chloride ion)
HCl (hydrogen chloride) in water (hydrochloric acid)
CN- (cyanide ion)
HCN (hydrogen cyanide) in water (hydrocyanic acid)
S2- (sulfide ion)
H2S (hydrogen sulfide) in water (hydrosulfuric acid)
compounds, d e.g., N NaCl Cl (sodium ( di chloride). hl id ) Given Formula: Name first as the element,
the second as an anion (root with –ide suffix).
Specify number of atoms of each element by using Greek prefixes (Table 2.4): Greek Prefixes Prefix MonoDiDi TriTetraPentaHexa-
Meaning 1 2 3 4 5 6
Prefix mono- never used for the first element.
When the prefix ends in a or name begins with a vowel, the prefix is often omitted. Formula CO CO2 PCl3 NO2 N2O5 SO3 SO2
o and the anion a or o of the
Name carbon monoxide carbon dioxide d d phosphorus trichloride nitrogen dioxide dinitrogen pentoxide sulfur trioxide sulfur dioxide
Hydrates Hydrate: A compound with a specific number of water molecules that are strongly adhered. CuSO4 Î copper(II) sulfate (anhydrous) C SO4 . 5H2O Î copper(II) CuSO (II) sulfate lf t pentahydrate t h d t BaCl2 Î barium chloride (anhydrous) BaCl2 . 2H2O Î barium chloride dihydrate Anhydrous: Dry, without water
What are the chemical names or formulas of the following compounds ? (NH4)2SO4 Hydroiodic acid (hydrogen iodide) KMnO4 Mercury(II) nitrite Fe(NO3)3 Potassium dichromate
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Note: An ion and the corresponding neutral species have completely different chemical properties. NO2
NO2-
SO3
SO32-
Na
Na+
Fe
Fe3+
H
H+
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