Atoms: The building Blocks of Matter
ATOMS: THE BUILDING BLOCKS OF MATTER Ms. Moore Chemistry
THE ATOM • Democrites- described natures particles as an atom( indivisible) • Three laws to describe behavior of matter • 1) The law of conservation of matter- states that matter can neither be created nor destroyed during chemical reactions or physical change. • 2) The law of definite proportions- a chemical compound contains the same elements in exactly the same proportions by mass regardless of the size of the sample or the source of the compound. • 3) The law of multiple proportions- if two or more different compounds are composed of the same two elements, then the ratio of the masses of the second element combines with a certain mass of the first element is always a ratio of small whole numbers.
LAW OF MULTIPLE PROPORTIONS
DALTON’S ATOMIC THEORY • Elements are made of small, indivisible particles called atoms • Atoms of an element are identical and have the same mass and properties • Atoms combine to form compounds in small, whole # ratios • Chemical Reactions are the rearrangement of atoms
STRUCTURE OF THE ATOM • Atoms contain positive and negative particles • JJ Thompson proved that electrons existed by using a cathode ray tube • 1) Cathode rays were deflected by a magnetic field in which a wire carrying electrical current is known to have a negative charge • 2) The rays were deflected away from a negatively charged object.
• He measured the ratio of the charge of cathode-ray particles to their mass. And found the ratio to be the same regardless of the metal. • -Named the particle electrons
• https://www.youtube.com/watch?v=2xKZRpAsWL8
• Milikans oil drop experiment- https://www.youtube.com/watch?v=XMfYHag7L
Millikan’s Oil Drop Experiment Showed that if a drop of oil were exposed to X- Rays, it became charged from the air By having the drop fall between two electrical plates
◦ Determined the average mass of a drop ◦ Drop was suspended in air using the electrical field ◦ By knowing what field was necessary to hold the drops, he was able to determine the charge
1.602 x 10-19 Coulombs
Media Resources\Rutherford Gold Foil.mov
• Rutherford’s Gold Foil Experiment
MODERN ATOMIC STRUCTURE • Nucleus • Protons – large, positively charged particle • Neutron – large, neutrally charged particle
• Electron – small, negatively charged particle Particle
Relative Mass
Relative Charge
Electron
1
-1
Proton
1836
+1
Neutron
1836
0
ISOTOPES • Atoms of an element must have the same # of protons • Atomic Number (Z) – number of protons
• Atoms of an element can have a different # of neutrons • Mass Number (A) – sum of protons and neutrons
• Atoms of an element can have a different # of electrons • Charge = (Protons – Electrons)
• Mass Number- is the total number of protons and neutrons that make up the nucleus of an isotope. • Mass number= Protons + neutrons • Atomic Number= number of protons= number of electrons. • Ex: how many protons electrons and neutrons are there in an atom of chlorine 37? • 1) The atomic number found on the periodic table is equal to the number of protons. • 2) The number of protons equals the number of electrons. • 3) The mass number=( protons + neutrons)- the atomic number equals the number of electrons.
PRACTICE • How many protons, electrons , and neutrons make up an atom of bromine80? • 35 protons 35 electrons and 45 neutrons
• Write the nuclear symbol for Carbon 13
• Write the notation for the isotope 15 electrons and 15 neutrons. • Phosphorous 30
CHARGED ISOTOPES • If the isotope is charged such as 37 Cl -1 add the charge to the number of protons to get a total number of electrons. •
P:
17 N: 20 E: 18 therefore making the isotope have a charge of negative 1
WEIGHTED ATOMIC MASS • The weighted average of the atomic masses of the naturally occurring isotopes of an element. • 1) convert the percentage to decimal form • 2) multiply the relative abundance by the atomic mass of each isotope. • 3) add up each result
• Ex: What is the atomic mass of Chlorine? • 35Cl: 75.77% 34.969 amu • 37Cl: 24.23% 36.966 amu • .7577 ∗ 34.969 + .2423 ∗ 36.966 = 35.46
ELECTRON CONFIGURATION
PERIODIC TABLE
PERIODIC TABLE • Developed by Mendeleev in 1869 • Placed elements in groups (columns) based on their properties
• Group (Family) • Vertical column of elements • Strongly related by properties
• Period (Row) • Related by a systematic change in properties
• Alkali Metals
GROUPS
• 1st column • Highly Reactive • Reactivity Increases down the table • Never found naturally in the pure state • Na found as Na+…K as K+ • Soft, Low Densities • React with H2O to make Base
2 Na( s) 2 H2 O(l ) 2 NaOH (aq) H2 ( g)
ALKALI METALS
• Alkaline Earth Metals
GROUPS
• 2nd Column • Reactive, but less than Alkali Metals • Reactivity increases down the Periodic Table • Never found naturally in pure state • Soft, but harder than Alkali Metals • React with H2O to make Base
Mg( s) 2 H2 O(l ) Mg(OH ) 2 (aq) H2 ( g)
ALKALI EARTH METALS
• Next to Last column
HALOGENS
• Highly reactive non-metals • Reactivity increases UP the Periodic Table • Rarely found naturally in the pure state • Powerful oxidizing agents
• Gases, Liquids and Solid • Colorful compounds in pure state
• Combine with Hydrogen to make Acids • HF, HCl, HBr, HI
H2 ( g) Cl2 ( g) 2 HCl ( g)
HALOGENS
NOBLE GASES • Last Column • Almost completely non-reactive gases • Very Few compounds exist • Melting points and Boiling points differ by less than 10 ºC • Argon is 1.3% of air (by mass) • Discovered and obtained from fractional distillation of air
NOBLE METALS • Very non-reactive metals • Highly resistant to corrosion or oxidation • Found naturally in the pure state • Copper, Silver, Gold • Clean surfaces in a vacuum stay clean • Many lists include other metals • Especially Platinum & Palladium • Both of these will have a dirty surface quickly (CO)
• Properties • Good conductors of heat and electricity • Shiny surface (high luster) • Malleable • Ductile • Found on the left side of the table • Solids at Room Temperature • Most elements on the Periodic Table
METALS
NON-METALS • Properties • Poor conductors (heat and electricity) • Colorful substances in the pure state
• Many gases and a few soft solids • Found on the extreme right side of the table
SEMIMETALS • Also called metalloids • Properties are between metals and non-metals • Stair-step of elements that separate the two main types • Everything to the left of the stair step • Metal
• Everything to the right of the stair step • Non-metal
OTHER GROUPS • Main Group • Elements in 1st two and last 6 columns • Make up the most abundant elements
• Transition Metals • Center Block of metals (called “d-block”)
• Inner Transition Metals • Block at the bottom of the table (“f-block”)
SUMMARY
Metals
Non-metals
NATURAL STATES • Most of the elements are found as a solid when pure (“standard state”) • Liquids • Mercury, Bromine • Gallium has a very low melting point (29.8 ºC)
• Gases • Noble Gases • Hydrogen, Nitrogen, Oxygen, Fluorine, Chlorine
ELEMENTS • Pure elements are not always individual atoms • Diatomic Elements – elements which are always found pure as a bonded pair of atoms • • • • • • •
Hydrogen (H2) Nitrogen (N2) Oxygen (O2) Fluorine (F2) Chlorine (Cl2) Bromine (Br2) Iodine (I2)
ALLOTROPES • Pure substances found naturally in different forms of the same physical state • Carbon – 40+ known allotropes • Graphite – chemically stable • Flat sheets of carbon
• Diamond – chemically unstable • 3-D bonded structure • Extremely hard and strong • 1 giant molecule
• Buckminster Fullerene • C60 – soccer ball structure (pentagon/hexagon)
THE ATOM • Democrites- described natures particles as an atom( indivisible) • Three laws to describe behavior of matter • 1) The law of conservation of matter- states that matter can neither be created nor destroyed during chemical reactions or physical change. • 2) The law of definite proportions- a chemical compound contains the same elements in exactly the same proportions by mass regardless of the size of the sample or the source of the compound. • 3) The law of multiple proportions- if two or more different compounds are composed of the same two elements, then the ratio of the masses of the second element combines with a certain mass of the first element is always a ratio of small whole numbers.
LAW OF MULTIPLE PROPORTIONS
DALTON’S ATOMIC THEORY • Elements are made of small, indivisible particles called atoms • Atoms of an element are identical and have the same mass and properties • Atoms combine to form compounds in small, whole # ratios • Chemical Reactions are the rearrangement of atoms
STRUCTURE OF THE ATOM • Atoms contain positive and negative particles • JJ Thompson proved that electrons existed by using a cathode ray tube • 1) Cathode rays were deflected by a magnetic field in which a wire carrying electrical current is known to have a negative charge • 2) The rays were deflected away from a negatively charged object.
• He measured the ratio of the charge of cathode-ray particles to their mass. And found the ratio to be the same regardless of the metal. • -Named the particle electrons
• https://www.youtube.com/watch?v=2xKZRpAsWL8
• Milikans oil drop experiment- https://www.youtube.com/watch?v=XMfYHag7L
Millikan’s Oil Drop Experiment Showed that if a drop of oil were exposed to X- Rays, it became charged from the air By having the drop fall between two electrical plates
◦ Determined the average mass of a drop ◦ Drop was suspended in air using the electrical field ◦ By knowing what field was necessary to hold the drops, he was able to determine the charge
1.602 x 10-19 Coulombs
Media Resources\Rutherford Gold Foil.mov
• Rutherford’s Gold Foil Experiment
MODERN ATOMIC STRUCTURE • Nucleus • Protons – large, positively charged particle • Neutron – large, neutrally charged particle
• Electron – small, negatively charged particle Particle
Relative Mass
Relative Charge
Electron
1
-1
Proton
1836
+1
Neutron
1836
0
ISOTOPES • Atoms of an element must have the same # of protons • Atomic Number (Z) – number of protons
• Atoms of an element can have a different # of neutrons • Mass Number (A) – sum of protons and neutrons
• Atoms of an element can have a different # of electrons • Charge = (Protons – Electrons)
• Mass Number- is the total number of protons and neutrons that make up the nucleus of an isotope. • Mass number= Protons + neutrons • Atomic Number= number of protons= number of electrons. • Ex: how many protons electrons and neutrons are there in an atom of chlorine 37? • 1) The atomic number found on the periodic table is equal to the number of protons. • 2) The number of protons equals the number of electrons. • 3) The mass number=( protons + neutrons)- the atomic number equals the number of electrons.
PRACTICE • How many protons, electrons , and neutrons make up an atom of bromine80? • 35 protons 35 electrons and 45 neutrons
• Write the nuclear symbol for Carbon 13
• Write the notation for the isotope 15 electrons and 15 neutrons. • Phosphorous 30
CHARGED ISOTOPES • If the isotope is charged such as 37 Cl -1 add the charge to the number of protons to get a total number of electrons. •
P:
17 N: 20 E: 18 therefore making the isotope have a charge of negative 1
WEIGHTED ATOMIC MASS • The weighted average of the atomic masses of the naturally occurring isotopes of an element. • 1) convert the percentage to decimal form • 2) multiply the relative abundance by the atomic mass of each isotope. • 3) add up each result
• Ex: What is the atomic mass of Chlorine? • 35Cl: 75.77% 34.969 amu • 37Cl: 24.23% 36.966 amu • .7577 ∗ 34.969 + .2423 ∗ 36.966 = 35.46
ELECTRON CONFIGURATION
PERIODIC TABLE
PERIODIC TABLE • Developed by Mendeleev in 1869 • Placed elements in groups (columns) based on their properties
• Group (Family) • Vertical column of elements • Strongly related by properties
• Period (Row) • Related by a systematic change in properties
• Alkali Metals
GROUPS
• 1st column • Highly Reactive • Reactivity Increases down the table • Never found naturally in the pure state • Na found as Na+…K as K+ • Soft, Low Densities • React with H2O to make Base
2 Na( s) 2 H2 O(l ) 2 NaOH (aq) H2 ( g)
ALKALI METALS
• Alkaline Earth Metals
GROUPS
• 2nd Column • Reactive, but less than Alkali Metals • Reactivity increases down the Periodic Table • Never found naturally in pure state • Soft, but harder than Alkali Metals • React with H2O to make Base
Mg( s) 2 H2 O(l ) Mg(OH ) 2 (aq) H2 ( g)
ALKALI EARTH METALS
• Next to Last column
HALOGENS
• Highly reactive non-metals • Reactivity increases UP the Periodic Table • Rarely found naturally in the pure state • Powerful oxidizing agents
• Gases, Liquids and Solid • Colorful compounds in pure state
• Combine with Hydrogen to make Acids • HF, HCl, HBr, HI
H2 ( g) Cl2 ( g) 2 HCl ( g)
HALOGENS
NOBLE GASES • Last Column • Almost completely non-reactive gases • Very Few compounds exist • Melting points and Boiling points differ by less than 10 ºC • Argon is 1.3% of air (by mass) • Discovered and obtained from fractional distillation of air
NOBLE METALS • Very non-reactive metals • Highly resistant to corrosion or oxidation • Found naturally in the pure state • Copper, Silver, Gold • Clean surfaces in a vacuum stay clean • Many lists include other metals • Especially Platinum & Palladium • Both of these will have a dirty surface quickly (CO)
• Properties • Good conductors of heat and electricity • Shiny surface (high luster) • Malleable • Ductile • Found on the left side of the table • Solids at Room Temperature • Most elements on the Periodic Table
METALS
NON-METALS • Properties • Poor conductors (heat and electricity) • Colorful substances in the pure state
• Many gases and a few soft solids • Found on the extreme right side of the table
SEMIMETALS • Also called metalloids • Properties are between metals and non-metals • Stair-step of elements that separate the two main types • Everything to the left of the stair step • Metal
• Everything to the right of the stair step • Non-metal
OTHER GROUPS • Main Group • Elements in 1st two and last 6 columns • Make up the most abundant elements
• Transition Metals • Center Block of metals (called “d-block”)
• Inner Transition Metals • Block at the bottom of the table (“f-block”)
SUMMARY
Metals
Non-metals
NATURAL STATES • Most of the elements are found as a solid when pure (“standard state”) • Liquids • Mercury, Bromine • Gallium has a very low melting point (29.8 ºC)
• Gases • Noble Gases • Hydrogen, Nitrogen, Oxygen, Fluorine, Chlorine
ELEMENTS • Pure elements are not always individual atoms • Diatomic Elements – elements which are always found pure as a bonded pair of atoms • • • • • • •
Hydrogen (H2) Nitrogen (N2) Oxygen (O2) Fluorine (F2) Chlorine (Cl2) Bromine (Br2) Iodine (I2)
ALLOTROPES • Pure substances found naturally in different forms of the same physical state • Carbon – 40+ known allotropes • Graphite – chemically stable • Flat sheets of carbon
• Diamond – chemically unstable • 3-D bonded structure • Extremely hard and strong • 1 giant molecule
• Buckminster Fullerene • C60 – soccer ball structure (pentagon/hexagon)
