Substances and their Properties Isotopes

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Substances and their Properties What are substances? - It can be identified as an element, compound, or a mixture. - A substance cannot be further broken down or purified by physical means. A substance is matter of a particular kind. Each substance has its own characteristic/ properties that are different from the set of properties of any other substance. Protons and Neutrons are found in the nucleus.

Element A chemical substance  Simplest substance  Built up from one type of atom only Atom – smallest particle of an element.    

Composed of: Protons – positively charged Neutrons - neutral Electrons – negatively charged

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Atomic number = number of protons = number of electrons Atomic mass = sum of protons and neutrons

Electrons circle the atoms in areas called shells.

Examples: Element Na

Atomic #

Atomic mass

Proton

Electron

Neutron

11

23

11

11

12

Ca

20

Ne

20

Al

20

10

13

14

Isotopes -

Isotopes are atoms of the same element (with the same number of protons) but different number of neutrons.

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Most elements that commonly occur are made up of isotopes. For example, chlorine exists as two isotopes. A sample of chlorine gas consists of - 75% of chlorine-35 and 25% of chlorine-37.

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Isotopes have the same chemical properties but slightly different physical properties. The chemical properties of isotopes are similar because chemical reactions involve only the electrons. The physical properties differ because the relative masses of the isotopes differ.

Atomic Structure Timeline Democritus (400 B.C.) • Proposed that matter was composed of tiny indivisible particles • Not based on experimental data • Greek: atomos John Dalton (1807)  British Schoolteacher  based his theory on others’ experimental data  Billiard Ball Model  atom is a uniform, solid sphere Dalton’s Four Postulates  1. Elements are composed of small indivisible particles called atoms.  2. Atoms of the same element are identical. Atoms of different elements are different.  3. Atoms of different elements combine together in simple proportions to create a compound.  4. In a chemical reaction, atoms are rearranged, but not changed. J. J. Thomson (1903)  Discovered Electrons  negative particles within the atom  Plum-pudding Model  positive sphere (pudding) with negative electrons (plums) dispersed throughout Ernest Rutherford (1911)  Discovered the nucleus  dense, positive charge in the center of the atom  Nuclear Model  Nuclear Model  dense, positive nucleus surrounded by negative electrons Niels Bohr (1913)  Energy Levels  electrons can only exist in specific energy states  Planetary Model  electrons move in circular orbits within specific energy levels

Erwin Schrödinger (1926)  Quantum mechanics  electrons can only exist in specified energy states  Electron cloud model  orbital: region around the nucleus where e- are likely to be found James Chadwick (1932)  Discovered neutrons  neutral particles in the nucleus of an atom

The Periodic Table of Elements     

The horizontal rows are called period and are labeled from 1 to 7. The vertical columns are called group/family are labeled from 1 to 18. An element is identified by its chemical symbol. The number above the symbol is the atomic number The number below the symbol is the rounded atomic weight of the element.

Dmitri Ivanovich Mendeleev: Father of the Table of Elements  The elements can be grouped into three broad classes based on their general properties.  Three classes of elements are Metals, Nonmetals, and Metalloids. Properties of Metals  Metals are good conductors of heat and electricity.  Metals are shiny.  Metals are ductile (can be stretched into thin wires).  Metals are malleable (can be pounded into thin sheets).  A chemical property of metal is its reaction with water which results in corrosion.  Solid at room temperature except Hg. Properties of Non-Metals  Non-metals are poor conductors of heat and electricity.  Non-metals are not ductile or malleable.  Solid non-metals are brittle and break easily.  They are dull.  Many non-metals are gases. Properties of Metalloids  Metalloids (metal-like) have properties of both metals and non-metals.  They are solids that can be shiny or dull.  They conduct heat and electricity better than non-metals but not as well as metals.  They are ductile and malleable. Boron, silicon, germanium, arsenic, antimony and tellurium.

Group I A – Alkali metals II A – Alkaline earth metals III A – Boron family IV A – Carbon family V A – Nitrogen family VI A – Oxygen family VII A – Halogen family VIII A or group zero – Noble gases

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Electron Configuration •

ELECTRONIC CONFIGURATIONS - is a way of distributing the electrons of the atom among the orbitals. Atomic Orbitals – are regions around the nucleus in which an electron stays most of the time. (s, p, d, and f) The sublevels of the electrons are divided into orbitals, with each orbital capable of accommodating a maximum of 2 electrons.

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Sublevel

Name of orbital

# of orbitals

Maximum # of Electrons

s

s orbital

1

2

p

p orbital

3

6

d

d orbital

5

10

f

f orbital

7

14

DETERMINING ELECTRON CONFIGURATIONS • FOR EXAMPLE: HYDROGEN (ATOMIC NUMBER =1) SO IT HAS ONE ELECTRON WHILE OXYGEN (ATOMIC NUMBER = 8) IT HAS EIGHT ELECTRONS. • BUT HOW ARE THEY ARRANGED WITH EACH ATOM? • IN ORDER TO WRITE A CORRECT CONFIGURATION FOR AN ATOM WE MUST: • (1) FIND THE NUMBER OF ELECTRONS THE ATOM CONTAINS • (2) USE THE MNEMONIC DEVICE

Ex. Oxygen- (atomic number = 8) electrons = 8 O= 1s22s22p4

Valence electrons and Lewis dot structure  Valence electron- Electrons found in the outermost energy of atoms, responsible for the reactions between elements.  How to determine the valence electrons? Example  Zn30 = 1s22s22p63s23p64s23d10  1= 2, 2= 8, 3=18, 4= 2 (Valence Electron)  Lewis dot structure(LEDS)- A convenient way of showing the arrangement of valence electrons around the atoms using dots as electrons and the symbol of the element as the nucleus. 

Electron arrangement determines the chemical properties of an atom • Electrons in an atom are arranged in shells, which may contain different numbers of electrons • Atoms whose shells are not full, tend to interact with other atoms and gain, lose, or share electrons. These interactions form chemical bonds. • The outermost electrons are the only ones involved in chemical bonding.

Types of chemical bonds: 1. Ionic bond 2. Covalent bond 3. Metallic bond WHY DO BONDS BETWEEN ATOMS FORM? • • •

WHEN BONDS FORM THE STABILITY OF THE COMBINED ATOMS INCREASES AS COMPARED TO THAT OF THE INDIVIDUAL ATOMS. GENERALLY CHEMICAL STABILITY IS RELATED TO THE ABILITY OF ATOMS TO ATTAIN THE ELECTRON CONFIGURATION OF AN INERT GAS. FOR MANY ATOMS THIS MEANS ACQUIRING EIGHT ELECTRONS IN THE OUTER SHELL. THIS IS CALLED AN “OCTET” STRUCTURE.

Ionic bonds are attractions between ions of opposite charge. (metal and nonmetal) • When atoms gain or lose electrons – Charged atoms called ions are created.

Covalent bonds join atoms into molecules through electron sharing. (Exists between nonmetals) • In covalent bonds, two atoms share one or more pairs of outer shell electrons, forming molecules. • Sharing of electrons may be EQUAL or UNEQUAL • A molecule is nonpolar when its covalently bonded atoms share electrons equally • In a polar molecule electrons are shared unequally between atoms, creating a polar covalent bond.

Chemical Reactions - Chemical reactions are processes in which substances (reactants) go through chemical changes to form new substances (products). - have two parts: - Reactants - the substances you start with - Products- the substances you end up with - The reactants turn into the products. - Reactants - Products

Rates of Reactions: In order for a reaction to occur: - The particles must collide (touch) - They must collide with enough energy - They must collide in the right orientation Factors that affect reaction rate: - Temperature (particle energy) - Particle size - Surface area - Particle contact (stirring) Evidence of Reactions: 1. Formation of a gas 2. Change in color 3. Formation of a solid (precipitate) 4. Change in heat or light energy Types of Chemical Reactions: 1. Synthesis- the get-together, Two or more chemicals bond together forming one new substance. Ex. 2Na + Cl2  2NaCl 2. Decomposition- the break-up Ex. 2H2O2  2H2O + O2 3. Single Replacement- the cheater, One element knocks another element out of a compound. Ex. 2HCl + Zn  ZnCl2 + H2 4. Double Replacement- the swap, Two compounds switch ions with each other. Ex. BaCl2 + Na2SO4  BaSO4 + 2NaCl 5. Combustion- Everyone loves O2 , A compound burns in oxygen gas. Ex. 2Mg + O2  2MgO

Balancing Chemical Equations LAW OF CONSERVATION OF MASS-Atoms can’t be created or destroyed - All the atoms we start with we must end up with - A balanced equation has the same number of each element on both sides of the equation. Rules for balancing:  Write the correct formulas for all the reactants and products  Count the number of atoms of each type appearing on both sides  Balance the elements one at a time by adding coefficients (the numbers in front)  Check to make sure it is balanced. Never - Change a subscript to balance an equation. - If you change the formula you are describing a different reaction. - H2O is a different compound than H2O2 - Never put a coefficient in the middle of a formula - 2 NaCl is okay, Na2Cl is not.