Introduction to atoms and molecules

Chapter 2-1 – 2-5

Chapter 5-7 and 5-9

Chapter 4-5 – 4-6

Key concepts in this unit

• Dalton’s atomic theory

• The electron, proton, and neutron

• Isotope, atomic number, mass number

• Molecules and molecular formulas

• Introduction to ions

• Nomenclature for inorganic compounds

Concept of the atom

• Democritus (ancient Greek) -- (atomos); indivisible particle of a substance

• Democritus’ ideas were promptly lost for the next ~ 2000 years or so.

Dalton—birth of the atomic theory (1803)

Dalton’s postulates:

1.Each element is composed of extremely small, indivisible particles called atoms.

2.All atoms of a given element are identical. Atoms of different elements are different, and have different properties, including mass.

Dalton’s atomic theory (con’t)

3. Atoms of an element are not changed into different types of atoms in a chemical reaction (i.e., you can’t turn lead into gold). Atoms are not created or destroyed in chemical reactions.

4. Compounds are formed when atoms of more than one element combine in small, whole number ratios.

5. A given compound always has the same relative number and kind of atoms.

• law of constant composition

• law of conservation of matter

• law of multiple proportions

Dalton’s theory explains laws

Law of multiple proportions

• Elements combining in more than one proportion will do so in multiples of whole numbers.

• Examples: – H2O vs. H2O2.

– NO NO2 N2O

Fundamental particles

• Masses of fundamental particles:

• Atomic mass unit: 1 amu = 1.66053873 10-24 g

Particle Mass (g) Mass (amu)

Proton 1.67262158 10-24 1.00727646

Neutron 1.67492716 10-24 1.00866492

Electron 9.10938188 10-28 5.485779911 10-4

Atoms

• Atoms vary by atomic number (Z), which indicates ________________________.

• Atoms contain the same number of electrons and protons (they have no charge).

• Atomic number increases from left to right on periodic table.

• The atomic weight of the atom is indicated below the symbol (in amu).

• Why are these weights not whole numbers?

Isotopes(chap 5-7, 5-9)

• Atoms of an element have same number of protons in the nucleus.

• Isotopes of an element vary in number of neutrons, but not in number of protons.

• Atomic symbols: (“carbon-12”)– C: carbon– Subscript number: atomic number; number of protons– Superscript number: mass number, number of

protons + neutrons

• What is the atomic symbol for carbon-14?

C126

Isotope symbols

Atomic symbol

Isotope name

Atomic number

Mass number

No. of protons

No. of neutrons

Chlorine-37

197 118

Na24_

Average atomic weight

• The average atomic weight (the one in the periodic table) is a weighted average.

• Not all isotopes of an atom exist in the same amounts, and this must be taken into account.

)_)(_( abundanceisotopemassisotopeAWavg

• Example 1: Let’s assume carbon has only 2 isotopes, 13C and 12C. The natural abundance of 13C is 1.0 %. What is the average atomic mass?

• Example 2: Cl has 2 isotopes, 35C and 37C. The average atomic mass of Cl is 35.5 amu. What is the natural abundance of the two isotopes?

Molecules

• A molecule is ___________________________________________________________________.

– In some cases, an atom is also a molecule. Example: He, Ne, Ar.

Some elements exist in nature as molecules.

– Diatomic molecular elements:

– Polyatomic molecular elements:

– Allotropes:

Formulas

• “shorthand” for writing and describing molecular composition.

– Chemical formula:

– Structural formula:

Ionic compounds

• Ionic compounds are composed of _________________________________.– Cation:– Anion:

• Ionic compounds do not form molecular units, but are _______________ containing a defined ratio of cations and anions, forming a _________.

• Formula unit:

Nomenclature for inorganic compounds

• The charge on an ion is designated by a number written above and to the right of the symbol

• Examples:– Na1+ (or Na+)– Al3+

– F-

– O2-

Periodic groups and ions

• Normally (but not always) atoms in the following groups form ions of the designated charge.

• This commonly occurs such that the # of electrons in the ion are equal to the # of electrons in the neighboring noble gas. We will discuss why this is so in quantum mechanics.

Group Common charge Alkali metals (1A) 1+ Alkaline earths (2A) 2+ Al or B 3+ Chalcogens (6A) 2- Halogens (7A) 1- Noble Gases (8A) 0

• Transition metals often form more than one type of cation (i.e., different charges, such as Fe+, Fe2+, Fe3+.)

• We have methods of writing formulas and names to indicate what charge the ions take.

Identifying ionic compounds

• Generally (but again, not always) ionic compounds are formed by combining a metal and a nonmetal. Molecular compounds are generally formed by combining nonmetals.

• The sum of the charges in an ionic compound must equal zero. The subscript of each ion indicates how many of that ion are in the formula.

• Example: what is the formula for the ionic compound formed when Al3+ combines with O2-?

Nomenclature rules

• Cations1. Cations from metal atoms take the name of the metal.

• Examples:2. Cations from transition metals have the charge indicated by a

Roman numeral.• Examples:– Note: there is an older method of naming transition metal cations.

We will not use the older method in this course, but you should be aware of it in case you ever run into it. (p.141)

3. Cations from nonmetal atoms end in –ium.• Examples:

Table 2-3 (p. 55) and Table 4-11 (p. 142) contain the names and symbols of common cations that you should know.

1. monoatomic anions are named by replacing the ending of the element with –ide.

– Examples:

Anions

2. polyatomic anions containing a variable number of oxygens. These occur in ternary salts.

The common anion is named using the suffix –ate.

NO3- nitrate

XO3- (X = F, Cl, Br, I) fluorate, chlorate, bromate, iodate.

CO32- carbonate

SO42- sulfate

PO43- phosphate

• The suffix or prefix of the name changes when the number of oxygens is increased or decreased

# of oxygens added or subtracted to common ion

prefix/suffix used

Add one oxygen per-___-ate common ion (+0 oxygens) -ate subtract one oxygen -ite subtract two oxygens hypo-__-ite

Note that the CHARGES on a “family” of oxyanions doesn’t change when thenumber of oxygens varies. The oxidation number of the central atom does change.We’ll discuss this later.

Examples of ternary salts

• sodium sulfate

• iron (II) oxide

• ammonium chloride

• lithium carbonite

– oxyanions with an H+ or 2 H+ added are indicated using the word hydrogen, or dihydrogen.

• examples:

(Sometimes the addition of a hydrogen is indicated by the prefix bi-. For instance, HCO3

- is the bicarbonate anion, as in NaHCO3, sodium bicarbonate. Use of either bi or hydrogen is correct).

Table 2-3 (p. 55) and Table 4-11 (p. 142) contain the names and symbols of common anions that you should know. You should also be able to properly name the relatives of the “common” oxyanions.

Acids

• For the moment, an acid is an ionic compound that, when dissolved in water, yields a H+.

anion prefix/suffix acid prefix/suffix example

-ide hydro-__-ic HCl, hydrochloric acid (from chloride ion)

per-___-ate -ate

per-___-ic -ic

HClO4, perchloric acid HClO3, chloric acid

-ite hypo-___-ite

-ous hypo-___-ous

HClO2, chlorous acid HClO, hypochlorous acid

Double check: the formula for any ionic compound must = ZERO.

Binary molecular compounds

1. The atom farthest to the left in the periodic table is usually written first.

2. For elements in the same group, the atom appearing lower in the table is usually written first.

3. the name of the second element in the formula name is given an –ide ending.

4. Greek prefixes indicate the number of each type of atom appearing in the molecule. (p. 142) You should know these prefixes.

one mono-

two di-

three tri-

four tetra-

five penta-

six hex-

seven hept-

eight oct-

nine non-

ten dec-