# Quantifying Chemical Compounds A YouTube element has been excluded from this version of the text. You can view it online here: http://pb.libretexts.org/chem12/?p=121

# The Law of Multiple Proportions

Elements combine in different ways to form substances whose mass ratios are small whole number multiples of each other. # The Law of Definite Proportions

#### All samples of a pure chemical substance always contain the same proportion of elements by mass regardless of their origin.

By mass, Carbon dioxide ( $CO_{2}$) is: 73% Oxygen and 27% Carbon

All carbon dioxide is chemically the same in every location.

# Return of the Mole!

##### We know:

One mole is equal to $6.022 \times 10^{23}$ objects.
One mole of any element equals its molar mass.

1 mole of carbon (C) has a mass of: 12.01 g Amedeo Avogadro – Inventor of the mole.

##### Molecules and the Mole:

One mole is equal to $6.022 \times 10^{23}$ objects.
One mole of any compound is equal to the sum of the molar masses of all elements in the compound.

1 mole of carbon dioxide $CO_{2}$ “>a mass of : $12.01g+16.00g+16.00g = 44.01g$ “>

Numbers of moles IS A DIRECT COMPARISON of the relative number of atoms or molecules!

The mole is used extensively in chemical problems because numbers of atoms can be compared! # Empirical Formula Vs. Molecular Formula

##### NOTE

The lowest whole integer numbers representing an atomic ratio of a molecule using a chemical description.

Empirical Formula of hydrogen peroxide: $HO$

A chemical description of the actual complete molecule.

Molecular Formula of hydrogen peroxide: $H_{2}O_{2}$ “>

Some compounds have empirical and molecular formulas that are the same.

Empirical Formula of water: $H_{2}O$
Molecular Formula of water: $H_{2}O$ “>

# Periodic Trends

##### Periodic Trends:
The periodic table is organized to help predict the properties of elements.
Elements down a column have similar chemical properties.
The periodic table is organized to help us determine useful information about elements. For example: atomic radius, electronegativity, and ionization energy. Learning periodic trends can help us understand why certain elements have the properties we observe.
##### Electronegativity The ability of a bonded atom to attract electrons Moving down a column on the periodic table electronegativity decreases Moving across a row on the periodic table electronegativity increases The difference in electronegativity determines bond type Electronegativity differences (ΔEN) for bonded atoms can be calculated by subtracting the least electronegative atom from the atom with the highest electronegativity. For hydrochloric acid (HCl): Electronegativity of Cl=3.0 Electronegativity of H=2.1 ∆����=��.��-��.��=��.��

 Is the bonding radius determined from averaging measurements of many compounds and molecules The bond length of a two bonded atoms is determined by adding their bond radii Moving down a column on the periodic table the principle quantum number (n) of the outermost electrons increases, orbitals are larger, therefore the atomic radii are larger. Moving across a row the electrons in the outermost shell feel the charge from the nucleus more strongly. Electrons are closer to the nucleus, orbitals are smaller, therefore the radii are smaller
##### On the periodic table bond radius increases down a column. On the periodic table bond radius decreases across a row. An atom that has lost an electron (cation) will have a smaller radius than the neutral atom. Fewer electrons = smaller electron cloud = smaller ionic radius. An atom that has gained an election (anion) will have a larger radius than the neutral atom. More electrons = larger electron cloud = larger ionic radius.

#### Cations have a smaller ionic radius than the neutral atom.

Anions have a larger ionic radius than the neutral atom. # Ionization Energy

 The amount of energy required to remove one electron from an atom (or ion) in a gaseous state. Moving down a column on the periodic table the principle quantum number (n) of the outermost electrons increase, orbitals are larger, therefore it takes less energy to remove an electron because they are farther from the nucleus. Moving across a row the electrons in the outermost shell feel the charge from the nucleus more strongly, orbitals are smaller, therefore it takes more energy to remove an electron because they are closer to the nucleus.
##### On the periodic table ionization energy decreases down a column.On the periodic table ionization energy increases across a row. ## Examples

###### Example 1

What is the molar mass of copper sulfate $CuSO_{4}$?

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###### Example 2

How many molecules of copper sulfate $(CuSO_{4})$ are in a 200 mg sample?

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###### Example 3

How many grams of copper are in a kilogram of copper sulfate $(CuSO_{4})$?

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###### Example 4

What is the empirical formula of a compound containing 94% oxygen and 6% hydrogen by mass?

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###### Example 5

A compound with the empirical formula $OH$ is found to have a molar mass of $34.0 \frac{g}{mol}$, what is the molecular formula?

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### Quiz

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# Chemical Bonding A YouTube element has been excluded from this version of the text. You can view it online here: http://pb.libretexts.org/chem12/?p=121

# Chemical Bonding Terminology  Two or more bonded atoms. Two or more bonded atoms of the same type.  Two or more bonded atoms of different types. Two or more molecules or atoms of the same type.

# Chemical Bonding

##### Polar Covalent

Electrostatic attraction between positive and negative ions.

Bonding atoms are generally metals and non-metals.

Electron is transferred from one element to another.

Electron sharing between atoms.

Bonding atoms are nonmetals.

Electrons are equally shared by atoms.

Electron sharing between atoms.

Bonding atoms are nonmetals.

Electrons are unevenly shared between atoms. Ionic Bonding in Sodium Chloride ( $NaCl$) Covalent Bonding in Hydrogen ( $H_2$) Polar Covalent Bonding in Hydrogen Fluoride ( $HF$)

# Ionic Bonding

##### Ionic bonding:

Metals have tendency to lose electrons, forming cations.

Nonmetals have a tendency to gain electrons, forming anions.

Metals and nonmetals want to attain the nearest noble gas configuration.

When metal and nonmetal atoms approach each other an electron transfer takes place.

The metal transfers its electron(s) to the nonmetal so both elements reach the nearest noble gas configuration.

The charged ions are then attracted to each other due to electrostatic forces.

#### Electrostatic Attraction Between Ions = Ionic Bond # Covalent Bonding

##### Covalent Bonding:
Atoms will try to attain a noble gas configuration.
Valence electrons are the electrons in the outermost shell of an atom.
Nonmetals will share valence electrons to satisfy a noble gas configuration.

#### Nonmetals Sharing Electrons = Covalent Bond # Polar Covalent Bonding

##### Polar Covalent Bonding:

An atom’s electronegativity is a measure of how well it attracts electrons to itself while in a chemical bond.

The difference in electronegativity between two bonded atoms can be calculated.

A covalent bond between atoms that have a significant difference in their electro-negativities.

Polar covalent bonds have unequal electron sharing between atoms.

#### Nonmetals Sharing Electrons Unequally = Polar Covalent Bond Molecules that have unequal electron sharing have partial charge separation. The atom in a molecule that is more electronegative will have a larger electron density and a slight net negative charge delta negative ( $\delta -$) The atom in a molecule that is less electronegative will have a smaller electron density and a slight net positive charge delta positive ( $\delta +$)

# Electronegativity and Bond Types

The electronegativity difference ( $\Delta EN$) between bonded atoms can be used to identify which type of bond exists in a molecule.

#### Electronegativity Differences and Bond Type

 ( $\Delta EN$) Bond Type Example 0.0-0.4 Covalent (no charge on atoms) $O_2$ Oxygen 0.4-2.0 Polar Covalent (partial charge on atoms) $CO$ Carbon Monoxide 2.0 + Ionic (full charge on ions) $KI$ Potassium Iodide # Valence Electrons and Bonding

##### Valence Electrons Continued:
Valence electrons are the electrons in the outermost shell of an atom.

The valence electrons are the electrons involved in chemical bonding between atoms.

The number of valence electrons an element has determines the chemical properties of that element.

Elements in a column (group) have the same number of valence electrons, this is why elements in a group have similar chemistry.

The valence electrons of a main group element are located in the outermost shell.

The valence electrons of a transition metal are located in the outermost d orbitals, as well as the outermost shell.

# Valence Electrons and Lewis Structures

##### Octet:
Representing the valence electrons of a main group element using dots surrounding the chemical symbol.

Chemical bonding is the attainment of a stable electron configuration through the sharing or transfer of electrons.

Lewis Theory uses the octet rule to predict bonding.

Lewis structures can be used to demonstrate bonding in molecules by showing atoms in a molecule sharing electrons to attain a full octet.

A full outer shell containing eight electrons.

# The Octet Rule

 A stable electron configuration can be attained with eight electrons in the outermost shell. Bonding atoms will transfer or share electrons to satisfy the octet rule, each atom will have access to eight electrons in its outermost shell. The octet rule can be used to predict how atoms bond. This rule only works for the second period (row) of the periodic table. Elements beyond the second row can access d or f orbitals, elements are larger and have more room for bonding. Hyper-coordination: Elements beyond the second row (after Neon) can have expanded octets. The octet rule is a useful tool for predicting bonding in molecules especially in organic chemistry.

# Lewis Structures Terminology

#### The charge on an atom in a Lewis structure, assuming all electrons are shared equally. Formal charge is calculated with the following formula: $(\text{valence }e^{-}) - (\frac{1}{2}\text{bonding }e^{-}) - (\text{lone electrons}) = \text{Formal Charge}$

# Lewis Structures

### First Row: ### Second Row and the Octet Rule: ### Lewis Structures and Bonding: # Lewis Structures for Molecules Method

##### Step 1:

Determine the total number of valence electrons for the entire molecule by adding the amount of valence electrons for each molecule, for anions add an extra electron, for cations subtract an electron.

##### Step 2:

Write the correct structure for the molecule, use a line to connect the bonding atoms, each line represents two electrons. Subtract bonding electrons from total electrons.

Hydrogen will always be terminal.

Generally the least electronegative atom will be in the center.

##### Step 3:

Distribute remaining electrons (as pairs) first to the terminal atoms, then to the central atom. Subtract from your total as you fill octets.

Hydrogen has a full shell with two electrons. Do not assign electrons (other than bonding) to hydrogen.

##### Step 4:

If there are any atoms without a full octet, move electrons to form double or triple bonds, use arrows to indicate electron movement. Redraw your structure.

## Examples

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### Quiz

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