Laws of Chemical Combination

Chem1 SY 10-11

-kaaferrer

What makes compounds different?

Law of Constant Composition 1799 Joseph Proust

a chemical compound contains the same elements in exactly the same proportions (ratios) by mass regardless of the size of the sample or source of the compound

For example, water always consists of oxygen and hydrogen atoms, and it is always 89 percent oxygen by mass and 11 percent hydrogen by mass

Does mass change during a chemical reaction?

Law of Conservation of Mass Lavoisier heated a measured amount of

mercury to form the red oxide of mercury. He measured the amount of oxygen removed from the jar and the amount of red oxide formed. When the reaction was reversed, he found the original amounts of mercury and oxygen.

Law of Conservation of Mass 1744 Antoine Lavoisier matter can not be created or destroyed in

ordinary chemical or physical changes. the mass of the reactants (starting materials)

equals the mass of the products

2Mg (s) + O2 (g) → 2MgO (s)

48.6 g 32.0 g 80.6 g

Example

10 grams of CaCO3 on heating gave 4.4g of CO2 and 5.6 of CaO. Show that these observations are in agreement with the law of conservation

Law of Multiple Proportions 1803 John Dalton

States that when two elements combine to form more than one compound, the masses of one element which combine with a fixed mass of the other element are in ratios of small whole numbers

Example: Carbon monoxide (CO): 12 parts by mass of

carbon combines with 16 parts by mass of oxygen.

Carbon dioxide (CO2): 12 parts by mass of carbon combines with 32 parts by mass of oxygen.

Ratio of the masses of oxygen that combines with a fixed mass of carbon (12 parts) 16: 32 or 1: 2

Water has an oxygen-to-hydrogen mass ratio of 7.9:1.

Hydrogen peroxide, another compound consisting of oxygen and hydrogen, has an oxygen-to-hydrogen mass ratio of 15.8:1.

Ratio of the masses of oxygen that combines with a fixed mass of hydrogen is 7.9: 15.8 or 1: 2

Chemical Formula Relationships

How many number of ATOMS? Ca(NO)3

NaNO3

Ba(IO3 )2

Molecular and Formula Weights The sum of all the atomic weights of each

atom in its chemical formula.

Formula mass = ∑ atomic masses in the formula unit

SEATWORK!

Percentage Composition from Formulas Percentage by mass contributed by each

element in the substance.

%element= (# of atoms of that element)(atomic mass of that element)

------------------------------------------------------------------------

formula mass of the compound

Stoichiometry

The quantitative relationships between the substances involved in a chemical reaction, established by the equation for the reaction

Terminologies

ATOM◦ Smallest particle of an element

MOLECULE◦ Smallest unit particle of a pure substance

ION◦ An atom or group of bonded atoms with electrical charge

because of an excess or deficiency of electrons

ELEMENT◦ Pure substance; CANNOT be broken down into 2 or more

pure substances by chemical means

COMPOUND◦ Pure substance; CAN be broken down into 2 or more pure

substances by chemical means

Atomic Mass

Atomic Mass = used to numerically indicate the mass of an atom in its ground state, it is expressed in the non SI unit of u

u = refers to unified atomic mass unit (formerly known as atomic mass unit or amu)

1 amu = 1/12 the mass of carbon-12 atom, therefore the mass of C-12 atom is made EQUAL to 12

amu  Carbon-12 atom is an isotope of carbon 

1 amu = 1.66 x 10-24 g

Atomic Mass

Mass of e = 1/1800 of mass of p and n so it is negligible making the equation

Atomic mass of 12C = mass of p + mass of n

subatomic particle

charge Mass (amu)

Neutron None 1.0087 ≈ 1

Proton Positive 1.0073 ≈ 1

Electron Negative 5.486 x 10-4 ≈ 0

Atomic Mass vs. Average Atomic Mass

Atomic Mass Average Atomic Mass

For carbon it is 12 u not 12.01 u

*Used to relate the fact that the numerical value assigned to each element in the periodic table reflects the average abundances of the atoms that compose a naturally occurring element

*Related to isotopes

*For carbon it is 12.01 u

*Chemists often will use the term “atomic mass” when they are actually referring to average atomic mass of an atom.

Average Atomic Mass (calculation)

Solve for the Average Atomic Mass of the element Boron

Average atomic mass = ∑ (mass x percent abundance)

Where ∑ means “sum”

Isotope Mass (u) Percent Abundance

11B 11.009305 80.110B 10.012937 19.9

Relative Atomic Mass & Atomic Weight

Have different meaning from atomic mass but synonymous with each other, although of different historical origins

Relative Atomic Mass “Ratio” of the average mass of the atom to the

unified atomic mass; dimensionless

Atomic Weight The name Dalton used in the early 19th century to

numerically describe the weight of atoms relative to each other

The average mass of the molecules of a binary compound (non metal-non metal)

Unit: u

Ex.: The molecular mass of carbon dioxide gas, CO2 is 28.01 u.

Molecular Mass

The average mass of the molecules of an ionic compound (metal-non metal)

Unit: u

Ex.: The formula mass of barium chloride, BaCl2 is 208.2 u.

Formula Mass

The Mole Concept

Used to describe the number of particles (atoms, molecules, etc.) that make up sample of matter

“One mole is the amount of any substance that contains the same number of units as the number of atoms in exactly 12 grams of carbon-12.”

1 mol of any substance = 6.02 x 1023 units of that substance,

where 6.02 x 1023 Avogadro’s Number, N

Molar Mass

Mass of 1 mole of substance SI unit: g/mol Provides a bridge between mass and amount

Conversion

GRAMS MOLESFORMULA

UNITS

Use molar mass

Use N

Atomic mass Molar mass

Unit: g/mol

Unit:amu; u

Atomic level

Macroscopic level

Numerically

equivalent

Comparison of Atomic and Molar Mass

Total massof p+ and no

Mass of 1 mole

Empirical Formula and Molecular Formula

Shows the relative number of atom of each element in the compound

EMPIRICAL FORMULA

A compound is found out to contain 20.0% carbon, 2.2% hydrogen, and 77.8% chlorine. Determine EF of the compound.

Problem 1

Shows the actual number of atom of each element in the compound

MOLECULAR FORMULA

MF = (EF)nwhere n is an integer

therefore,

n = MMF

MEFwhere M is molar mass

Relationship between EF & MF

A compound with the empirical formula C2H5 has a molar mass of 58.12 g/mol. Find the molecular formula of the compound.

Problem 2

The molar mass of a compound that is 54.6% carbon, 9.0% hydrogen and 36.4% oxygen is 88 g/mol. Find MF of the compound.

Problem 3

The Mole

One mole is the amount of any substance that contains the same number of units as the number of atoms in exactly 12 grams of carbon-12.

6.02 x 1023 Avogadro’s Number

MOLAR Mass The mass in grams of one (1) mole of a

substance. g/mol, grams per mole

The molar mass of any substance in grams per mole is always numerically equal to the atomic, formula, or molecular mass of the substance in amu.

moleatom

s gramFormula weight

Avogadro’s

number