Chapter 2

Atoms

• Matter is anything that has mass and occupies space.

• Made up of atoms• Examples of things that are not matter include

thoughts and emotions.• Classified as being either a pure substance or a

mixture of two or more pure substances.

Matter

• Pure substances are defined as having a “fixed composition” - one exact combination of ingredients that cannot be varied.

• Ingredients of a pure substance are chemically combined so that they lose their original identity and cannot be separated out without major chemical intervention.

Pure Substances

• Elements are substances which consist of only one type of atom. All the elements are listed on the periodic. Oxygen

• Compounds are substances which consist of more than one type of atom combined in a fixed ratio. Water

Pure Substances

Elements

• The “building blocks of nature”. All substances, all matter is created from elements.

• Elements are given chemical symbols, which are one or two letter abbreviations derived from the common or Latin name of the element.

• Chemical symbols are always written with the first letter capitalized and the second letter (if there is one) in lower case.

• Mixtures are two or more pure substances mixed in a “varying ratio”. This means that there is more than one way to combine the ingredients.

• In addition, the ingredients are physically mixed but not chemically combined and they can be easily separated.

• Trail mix, hot chocolate

Mixtures

• Homogeneous mixtures (solutions) have a uniform (same) appearance.

• The ingredients are evenly mixed and typically you cannot see the individual ingredients.– Kool aid

• If you take two samples of a homogeneous mixture, the samples will have the same composition due to the even mixing. – Glasses of Kool aid

Homogenous Mixtures

• Heterogeneous mixtures do not have a uniform (different) appearance.

• The ingredients do not mix evenly and typically you can see the individual components. – Trail mix, Chocolate chip cookies.

• Samples of heterogeneous mixtures do not necessarily have the same composition. – Hand fulls of trail mix

Heterogeneous mixtures

Classifying Matter

Oxygen Water Kool Aid Trail Mix

• Classify the following substances:• Ice

• Blueberry muffin

• Vitamin A

• Oxygen

• Tea

• Vegetable soup

Classifying Matter

John Dalton (1766-1844)

• English chemist & teacher (started teaching at age 12)

• Experiments led to discovery that the ratios of the

elements in the compound are whole number multiples

(law of multiple proportions)

Dalton’s Atomic Theory (1808)• First modern atomic theory, based on empirical evidence• Called particles atoms in honor of Democritus• Based on 3 important laws Law of Conservation of Matter

Law of Constant Composition Law of Multiple Proportions• Dalton proposed an atomic theory for matter that would

explain the existence of all 3 laws.

Postulates of Dalton’s Atomic Theory

1. All elements are made of tiny, indivisible particles called atoms.

2. All atoms of a given element have the same chemical properties. Conversely, atoms of different elements have different chemical properties.

3. In ordinary chemical reactions, no atom of any element disappears or is changed into an atom of another element

4. Compounds are formed by the chemical combination of two or more different kinds of atoms. Ina given compound, the relative numbers of atoms of each kind of element are constant, and are most commonly expressed as integers.

5. A molecule is a tightly bound combination of two or more atoms that acts as a single unit.

Law of Conservation of Mass

During a chemical reaction, the mass of the reactants before the reaction is _____ to the mass of the products after the reaction (mass is “conserved”)

Mass of reactant100.00g

Mass of products = 93.57 g + 6.43 g 100.00 g

Law of Constant Composition

Wherever found, calcium carbonate is always:40.04% calcium12.00% carbon47.96% oxygen

Monatomic Elements

• Consist of single atoms that are not bonded to each other.

• Helium• Krypton• Neon

Diatomic Elements• Contain 2 atoms bonded to each other in a

molecule. Usually exist as gasses and will not exist as “free atoms”

• Oxygen gas: O2

• Hydrogen gas:• Nitrogen gas:• Fluorine gas:• Chlorine gas:• Bromine gas: • Iodine gas:

Polyatomic Elements

• Contain many of the same atoms bonded to each other

• Ozone• Figure 2.5 page 37

The Atom

• Made up of 3 subatomic particles– Proton– Electron– Neutron

Subatomic Particles

Particle Location

Relative electrical charge

Actual

Mass (g)

Electron In space around the nucleus

9.11 x 10-28

Proton In the

nucleus

1.673 x 10-24

Neutron In the

nucleus

1.675 x 10-24

RelativeMass (amu)

__1__

1840

1.0

1.0

Rutherford’s Nuclear Model of the Atom

• The atom is mostly empty space (99.99+%)

• The positive charge is centrally located in a dense nucleus (most of the mass of the atom)

• The nucleus is approximately 100,000 times smaller than the atom

• The electrons are located at a distance away from the nucleus and avoid being pulled into the nucleus by orbiting at rapid speeds

Rutherford’s Atomic Model(“planetary model”)

e-

__________Diameter = 10-15 m

_________________

Diameter of atom = 10-10 m

• Carbon-12– ___ protons– ___ neutrons– ___ electrons

Mass Number

• The mass number is the number of _____ and ______ in an atom.

• The mass of an atom comes from the ______ and _______, because the _______ had “negligible” mass.

• Problem 2.2 page 39

Atomic Number (Z)

• The atomic number is the number of __________ in an atom.

• It is unique to each type of atom. For example, all carbon atoms have ____ protons. Any atom with six protons is a ______ atom.

• The atomic number also gives us the number of _______ in an atom.

• Atoms are charge-neutral, which means that every proton in the nucleus is balanced by an electron in the electron cloud.

Atomic Number

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Carbon has an atomic number of __

Chlorine has an atomic number of __

Sodium has an atomic number of __

Problem 2.3 and 2.4 page 40

Element Symbol Atomic Number

Mass Number

Protons Neutrons Electrons

N 15

Calcium 42

38 50

14 16

56 138

Problem

What is an isotope?

A= mass number Z = atomic number

Isotopes differ only in numbers of neutrons

Atoms with the same number of protons but different numbers of neutrons are called isotopes.

Isoptopes

• It is possible for atoms of the same element to have different numbers of neutrons.

• The identity of the element is based on the number of protons

• The neutrons are charge-neutral, changing the number of neutrons only changes the mass of the atom.

Measuring atomic masses

Isotope % found196Hg 0.15198Hg 9.97199Hg 16.87200Hg 23.10201Hg 13.18202Hg 29.86204Hg 6.87

Magnesium in Nature

Magnesium has three naturally-occurring isotopes:24Mg 78.99 %25Mg 10.00 %26Mg 11.01 %

1. How many protons, neutrons & electrons in each isotope?

2. What is the principal isotope of Mg

Bromine in Nature

Bromine has two naturally-occurring isotopes:79Br 50.69 %81Br 49.31 %

1. How many protons & neutrons in each?

2. In 10,000 atoms of Br, how much of each isotope would be present?

Problem 2.5 page 41

Atomic Mass

• The mass of an atom is based on the average mass of all the isotopes of an element.

• Elements have a naturally occurring abundance of each isotope.

• These naturally occurring abundances are used with the isotopic mass to calculate the average atomic mass of an element.

Atomic Mass Calculations

• The atomic masses of each element are shown on the periodic table.

• When you do these problems, you can always check your work against the periodic table.

• Your results should be close but will probably not be exactly the same.

Magnesium has three naturally-occurring isotopes:24Mg 78.99 %25Mg 10.00 %26Mg 11.01 %

1. What is the average atomic mass of Mg?

Bromine has two naturally-occurring isotopes:79Br 50.69 %81Br 49.31 %

2. What is the average atomic mass of Mg?

Atomic Mass Calculations

• Copper has two isotopes: 63Cu with a mass of 63.01 amu and an abundance of 69% and 65Cu with a mass of 65.01 amu and an abundance of 31%. Calculate the atomic mass of copper.

• In 100 g of copper, how many grams of 63Cu?

Dmitri Mendeleev (1834-1907)• Chemists began trying to design

a system of classifying the elements based on their properties

• When Mendeleev arranged the elements by increasing atomic weight, he noticed a periodic repetition in atomic properties (e.g. density, melting point, etc.)

• Mendeleev left blanks in his table, predicting elements (and their properties) not yet discovered

• Led to search for these elements that, when discovered, had the properties Mendeleev had predicted

• Powerful predictive properties of the table led to its acceptance

Mendeleev’s Table

The Periodic Table

• The periodic table is a systematic arrangement of all the elements which exist in:– Nature

– contrived laboratory conditions

– In our minds (?)

The Periodic Table

• Groups: Vertical columns

• Periods: Horizontal rows

• Elements in the same groups have similar chemical properties.

The Periodic Table

• Metals: To ______ of jagged line (except hydrogen)• Nonmetals: To _____ of jagged line (except hydrogen)• Metalloids: _______ jagged line (except aluminum

which is a metal)• Alkali metals: Group ____• Alkaline earth metals: Group ____• Halogens: Group ____• Noble gases: Group ____• Transition metals: All Group ____

Metals, Nonmetals, and Metalloids

• Metals are shiny, ductile and malleable, and are good conductors of heat and electricity. They have _____ densities and _____ melting points.

• Nonmetals are dull and brittle. They are poor conductors of heat and electricity and have relatively _____ densities and _____ melting points.

• Metalloids are shiny and brittle. They are poor conductors of heat and electricity, and have _____ densities and melting points (although not as _____ as metals).

Using the Periodic Table

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Electron Energy Levels

• Electrons are constantly moving• The shape of the electron path is called the orbital

– Sphere shapes are called s orbitals– Figure 8 shapes are called p orbitals

• The distance between electrons and the nucleus is called the energy level– Closest to the nucleus is called n=1– Energy levels go to infinity

• Practice worksheet from your Chem 121 packet

• Electrons are free to move between energy levels, like going up and down floors in a building– Moving up in energy levels requires atom to absorb energy– Moving down in energy levels requires atom to release

energy• Energy absorbed and released during energy level

transitions is called photons• Photons are more commonly known as “light” and

“radiation”

Electron Energy Levels

Valence Electrons

• We now know that periodic trends are the result of similarities in the electrons in an atom. More specifically, periodic trends are the results of similarity in valence electrons, which are the electrons located in the highest energy level (furthest from the nucleus).

• Valence electrons perform the chemistry in a chemical reaction

Valence Electrons

• All atoms in Group 1A have one valence electron, Group 2A have 2 valence electrons, Group 3A have 3 valence electrons etc…

• Valence electrons are typically represented in one of two ways: electron configuration (not covered in this class) and Lewis dot symbols (also called electron-dot symbols).

Lewis Dot Symbols (Structures)

• Lewis dot symbols show only the valence electrons in an atom.

Periodic Trends

• The periodic table arranges elements according to similarities in properties, such as reactivity or physical state.

• These similarities are called periodic trends.

Periodic Trends

• The periodic table was originally designed based completely on observations of periodic trends.

• The chemists who created the periodic table were even able to correctly predict physical properties and the reactivity of elements that had not yet been discovered.

Periodic Properties

• Atomic Radius is the volume of the electron clouds in atoms

• Increases going down each group• Decrease going across from left to right

Atomic Radius

Increasing AtomicRadius

DecreasingAtomic Radius

Periodic Properties

• Ionization Energy is the energy needed to remove the least tightly bound electron from an atom in the gaseous state. When this happens, the atom becomes an ion

• Increases going across from left to right• Decreases going down a group

Ionization Energy

Ionization Energy

Decreases

Ionization Energy Increases