Atoms and Molecules

By Doba Jackson, Ph.D.

Outline of Concepts in this chapter

• Law of Constant Composition (sec. 2-4)

– Mass Percentages• Dalton’s Atomic Theory (sec. 2-5)

– Postulates• Atomic and Molecular Mass (sec. 2-7, 2-8)

– Atomic Mass Unit (amu)• Protons, Neutrons, and Electrons (sec 2-9)

Dalton’s Atomic Theory

1. Elements are made up of tiny particles called atoms.

2. Each element is characterized by the mass of its atoms.

3. Chemical combination of elements to make different chemical compounds called molecules.

4. Chemical reactions only rearrange the way that atoms are combined in molecules; the atoms themselves don’t change.

Law of Constant Composition

2Hg2HgO

Chemical Formula

O2+

Chemical Equation

1) All matter is made of a specific composition of elements from the periodic table.

2) The mass percentage of each element in a pure molecule is conserved.

Law of Constant Composition

Chemical Formula

2Hg2HgO O2+

Chemical Equation

HgO

Mass percentage of mercury

%4.7100

%6.92100

xHgOofmass

oxygenofmass

xHgOofmass

mercuryofmass

Mass percentage of oxygen

Suppose we analyze 2.83 g of Lead Sulfide, PbS (consists of lead and sulfur), and find that it consist of 2.45 g of lead (Pb) and .380 g of sulfur (S). Calculate the mass percentages of lead (Pb) and sulfur (S) in the sample.

Example

%4.1310083.2

380.100

%6.8610083.2

45.2100

xg

gx

PbSofmass

Sofmass

xg

gx

PbSofmass

Pbofmass% Lead (Pb)

% Sulfur (S)

Definitions that lead to Dalton’s Atomic Theory

Law of Constant Composition: The relative amount of each element in a particular compound is always the same regardless of the source of the compound, or how the compound was prepared.

Law of Multiple Proportions: This law says that when a given element, (X), combines with another element, (Y), then the ratio of masses of X and Y are small whole numbers.

Law of Conservation of Mass: For any closed system, the mass of the system will remain constant over time

Law of Multiple Proportions

Law of Multiple Proportions: Elements can combine in different ways to form different compounds, with mass ratios that are small whole-number multiples of each other.

CompoundGrams of oxygen

Grams of nitrogen

Ratio

Nitric Oxide

8.10 grams 7.60 grams 1.07 : 1.00 NO

Nitrous Oxide

16.3 grams 8.1 grams 2.01 : 1.00 NO2

Law of Multiple Proportions

Law of Multiple Proportions: Elements can combine in different ways to form different compounds, with mass ratios that are small whole-number multiples of each other.

Law of Conservation of Mass: Mass is neither created nor destroyed in chemical reactions.

Aqueous solutions of mercury(II) nitrate and potassium iodide will react to form a precipitate of mercury(II) iodide and aqueous potassium iodide.

Law of Conservation of Mass: Mass is neither created nor destroyed in chemical reactions.

HgI2(s) + 2KNO3(aq)Hg(NO3)2(aq) + 2KI(aq)

4.55 g + 2.02 g = 6.57 g

3.25 g + 3.32 g = 6.57 g

Dalton’s Atomic Theory

1. Elements are made up of tiny particles called atoms.

2. Each element is characterized by the mass of its atoms.

3. Chemical combination of elements to make different chemical compounds called molecules.

4. Chemical reactions only rearrange the way that atoms are combined in molecules; the atoms themselves don’t change.

Molecules are groups of atoms joined together

Each element has a unique individual mass (atomic mass)

Atomic mass: The mass of a single element

Molecular mass: The mass of a single molecule

Atomic masses can be used to find mass percentages

Find the mass percentages of lead (Pb) and sulfur (S) in lead sulfide (PbS). Use theperiodic table.

Molecular mass of PbS

Atomic mass of Pb

Atomic mass of S =+

207.2 amu + 32.07 amu = 239.3 amu

% Lead (Pb)

%4.131003.239

07.32100

%6.861003.239

2.207100

xg

gx

PbSofmass

Sofmass

xg

gx

PbSofmass

Pbofmass

% Sulfur (S)

What is the atom made of?

Cathode-Ray Tubes: J. J. Thomson (1856–1940) proposed that cathode rays must consist of tiny, negatively charged particles which we now call electrons.

JJ Thompson reasoned that these cathode rays come from negatively charged particle called electrons.

Subatomic particles of the atom

Atomic Structure: Protons and Neutrons

Rutherford’s Scattering Experiment

Ernest Rutherford (1871–1937) directed a beam of alpha particles at a thin gold foil.

A very small number, (about 1 in every 20,000) were deflected at an angle.

Rutherford’s Scattering Experiment

Rutherford proposing that an atom must be almost entirely empty space and have its mass concentrated in a tiny central core that he called the nucleus.

A comparison of subatomic particles

The mass of the atom is primarily in the nucleus.

A comparison of subatomic particles

The charge on a proton is opposite the charge of the electron

Atomic Numbers, Mass Numbers, Isotopes

Atomic Number (Z): Number of protons in an atom’s nucleus. Equivalent to the number of electrons around the atom’s nucleus.

Mass Number (A): The sum of the number of protons and neutrons in an atom’s nucleus. This is the atomic mass in units of amu (1 amu = 1.661 x 10-24 g).

Isotopes: Atoms with identical atomic numbers but different mass numbers.

How to write an atoms atomic numbers, mass numbers for different Isotopes

Carbon-14

C14

6

Atomic number

Mass number

Carbon-12

C12

6

Atomic number

Mass number

6 protons6 electrons8 neutrons

6 protons6 electrons6 neutrons

Atomic Mass, Atomic Number

The mass of 1 atom of carbon-12 is defined to be 12 amu.

Atomic Mass: The weighted average of the isotopic masses of the element’s naturally occurring isotopes.

Why is the atomic mass of the element carbon 12.01 amu?

Carbon-12: 98.89 % natural abundance 12.0000 amu

Carbon-13: 1.11 % natural abundance 13.0034 amu

Carbon-14: 1 x10-12 % natural abundance 14.0032 amu

= 12.01 amu

= (12.00 amu)(0.9889) + (13.0034 amu)(0.0111)

= 11.87 amu + 0.144 amu

Average mass of a Carbon atom

Atomic masses on the periodic table are weighted averages of all naturally occurring isotopes