Lecture 02 (Chapter 2) Atoms, Molecules, and Ions

Daniel L. RegerScott R. GoodeDavid W. Ball

http://academic.cengage.com/chemistry/reger

Lecture 02 (Chapter 2)Atoms, Molecules, and Ions

• John Dalton (1766-1844) used experimental data to propose one of the first theories to explain the properties of matter.

• Dalton’s Atomic Theory, based on the idea that matter is discontinuous, is described by 4 postulates.

• Dalton’s Atomic Theory helped to explain other observations, including the law of constant composition.

Dalton’s Atomic Theory

1. Matter is composed of atoms. An atom is the smallest unit of an element that has all the properties of that element.

2. An element is composed entirely of one type of atom.

3. A compound contains atoms of two or more different elements. The relative number of atoms of each element in a compound is always the same.

4. Atoms do not change identity in chemical reactions; only the way in which they are joined together changes.

Dalton’s Atomic Theory

• Law of constant composition: All samples of a pure substance contain the same elements in the same proportions by mass.• From Dalton’s third assumption.

• Law of multiple proportions: When the same elements form more than one compound, the masses of one element that combines with a fixed mass of a second element are in a ratio of small whole numbers.• From Dalton’s third assumption.

• Law of Conservation of Mass: There is no detectable change in mass when a chemical reaction occurs.• From Dalton’s fourth assumption.

Laws Related to Dalton’s Atomic Theory

• Experiments over many years showed that atoms are not simple particles, but are composed of the subatomic particles listed below:• Electrons• Protons• Neutrons

Atomic Composition and Structure

Mass (g) Mass (u)

Proton (p+) 1.67 x 10-24 1

Neutron (n) 1.67 x 10-24 1

Electron (e-) 9.07 x 10-28 1/1836

• Experiments in the late 1800’s found that the application of high voltage across a partially evacuated tube produced cathode rays.

Cathode Rays; Discovery of electrons

• In 1897, J. J. Thomson demonstrated that cathode rays were negatively charged by applying magnetic and electric fields to cathode rays, which deflected their path, and confirmed their negative charge.

• Cathode rays are electrons, negatively charged particles that are one of the components of an atom.

Electrons

• 1913, Robert Millikan performed experiments to measure charge of electron.• Oil drops formed by injector were charged by capturing electrons. Based on

rate the drops moved in absence or presence of electrical field, charge of the electron was calculated to be 1.602 x 10-19 coulombs.

Millikan Oil Drop Experiment

• 1899, Ernest Rutherford conducted experiments that involved passing alpha particles (from radioactive decay) through thin metal targets.

• Observed that some particles deflected at very large angles.

The Nuclear Model of the Atom

• Rutherford concluded that the results of the scattering experiment required that atoms consist of:• a nucleus that is very

small compared to the atom, has a high positive charge and contains most of the mass of the atom.

• the remainder of the space in an atom contains enough electrons to give a neutral atom.

The Nuclear Model of the Atom

• Rutherford proposed that the hydrogen nucleus was a fundamental particle called the proton, which has a positive charge equal in magnitude to the negative charge of the electron.• Protons account for charge on nuclei of all atoms.• Proton mass (1.673 x 10-27 kg) is 1836 times that of the

electron.• However, the number of protons in a nucleus accounted for

half or less of the nuclear mass.• Scientists inferred there must be a massive, neutral

particle also present in the nucleus.• This neutral particle is called the neutron; its mass is almost

the same as that of the proton.

Protons and Neutrons

Particle Charge (C) Mass (kg) Relative charge

Relative mass

Electron -1.602 x 10-19 9.109 x 10-31 1- 0

Proton +1.602 x 10-19 1.673 x 10-27 1+ 1

Neutron 0 1.675 x 10-27 0 1

Particles in the Atom

• Atomic number (Z) is the number of protons in the nucleus of an atom (see periodic table).

• Mass number (A) is the sum of the numbers of protons and neutrons in the nucleus (may indicate different isotopes, i.e., same Z/different A of same element).

Definitions

XAZH1

1

H21

H31

Applications of Atomic and Mass Numbers

• On the periodic table, the atomic number is written as a whole number above the symbol F.

• In the written description, fluorine is said to have 9 protons (the atomic number is the number of protons).

• In the symbol, the number 9 is written in the atomic number or Z (lower left) position.

• Note: The periodic table does not show the mass number for an individual atom. It lists an average mass number for a collection of atoms!

Seager SL, Slabaugh MR, Chemistry for Today: General, Organic and Biochemistry, 7th Edition, 2011

Isotopes• Isotopes are atoms that have the same number of

protons in the nucleus but different numbers of neutrons. That is, they have the same atomic number but different mass numbers.

• Because they have the same number of protons in the nucleus, all isotopes of the same element have the same number of electrons outside the nucleus.


Symbols of Isotopes

• Oxygen also has three isotopes, containing 8, 9, and 10 neutrons respectively. The symbols are:

• Since the value of Z, and the symbol, both identify the element, Z is often omitted from the symbol:

O O O 188

178

168

O O O 181716

Elemental Notation

Pb20882

Q: How to represent lead-208?

Q: How many p+, e-, n? 82, 82, 126

X94

Q: How to represent element X with 4 p+ and 5 n.

• In many chemical reactions, atoms gain or lose electrons, producing charged particles called ions.• A cation has a positive charge and forms when an atom loses one or

more electrons (Na+, Pb2+).• An anion has a negative charge and forms when an atom gains one

or more electrons (Cl-, O2-).• In many cases, the periodic table tells us whether an atom tends to

lose or gain electrons, and how many.

Common atomic ions you should know:• H+, Na+, K+, Mg2+, Ca2+, Fe2+, Fe3+, Ag1+, Pb2+, • N3-, P3-, O2-, S2-, F-, Cl-, Br-

Ions

• The atomic number and atomic mass for each element is given on the periodic table.

Sr38

87.62

Atomic number

Atomic mass

Atomic Number and Atomic Mass

• Write the symbols for the particles containing:(a) 8 protons, 9 neutrons, 10 electrons(b) 13 protons, 14 neutrons, 13 electrons

Test Your Skill

Test Your Skill

• Write the symbols for the particles containing:(a) 8 protons, 9 neutrons, 10 electrons(b) 13 protons, 14 neutrons, 13 electrons

Answer: (a) (b)2178O Al27

13

Example: Components of Ions

• Fill in the blanks.Symbol

Atomic number____Mass number ____Charge ____no. of protons ____no. of neutrons____no. of electrons ____

Na2311

• A relative mass scale has been established to express the masses of atoms.

• The atomic mass unit (u) is 1/12 the mass of one 12C atom. Experimentally to three significant digits:

1 u = 1.66 x 10-27 kg• The masses of both the proton and the neutron are

approximately 1 u.• A 24Mg atom has a mass approximately twice that of

the 12C atom, so its mass is 24 u.• A 4He atom has a mass approximately 1/3 that of the

12C atom, so its mass is 4 u.

The Atomic Mass Unit (u)

• Factors other than the mass of the protons and neutrons affect the mass of atoms, so the actual mass of atoms are not whole numbers. (24Mg = 23.98504 u; 4He = 4.002603 u)

• When the accurate atomic mass of an atom is rounded to a whole number, it equals the mass number.

• Why are the numbers different?

Atomic Mass and Mass Number

How Isotopes Determine Atomic Mass

• The atomic mass of an element is the relative mass of an average atom of the element expressed in atomic mass units.

• Many elements have more than 1 isotope (e.g. – 12C, 13C, 14C). • Abundance of isotopes are not evenly distributed. • Weighted atomic mass of Carbon (12C, 13C only) = (0.98882*12u)

+ (0.01108 * 13.300335u) = 12.011u.

C612.011

12

C612.011

12C12

6 C13

6 C14

6 Carbon-12 Carbon-13 Carbon-14

%Abundance 98.892 1.108 1.0 x 10-10 AMU 12 u 13.300335 u

100

mass isotope% isotope mass Atomic

• A mass spectrometer is used to measure the relative masses of isotopes, neon in this figure.

The Mass Spectrometer

The mass spectrum of neon shows that the element is a mixture of three isotopes.

Mass Spectrum of Neon

• About 75% of the elements occur in nature as mixtures of isotopes.

• Usually, the relative abundance of isotopes of an element is the same throughout nature.• In all natural samples of Li, 7.42% of the atoms are 6Li

and the remaining 92.58% are 7Li.

Natural Distribution of Isotopes

Determining Atomic Mass

• A specific example of the use of the equation is shown below for the element boron that consists of 19.78% boron-10 with a mass of 10.01 u and 80.22% boron-11 with a mass of 11.01u.

• This calculated value is seen to agree with the value given in the periodic table.

u 81.10100

u 883.2 u 198.0

100

)u 01.11%22.80u 10.0119.78% AM


• A mass spectrometer was used to determine that gallium is 60.11% 69Ga (isotopic mass = 68.9256 u) and 39.89% 71Ga (isotopic mass = 70.9247 u). Calculate the atomic mass of Ga.

Example: Calculating Atomic Mass

• Proposed independently by Dimitri Mendeleev and Lothar Meyer. • Periodic table: arranges the elements in rows that place elements

with similar properties in the same column.• Period: a horizontal row• Group: a column - contains chemically similar elements

The Periodic Table

http://web.sbu.edu/chemistry/wier/atoms/meyer.html; http://www.chemistry.co.nz/mendeleev.htm

The Periodic Table

• Metal: a material that is shiny and is a good electrical conductor; metallic elements are on the center and left side of the periodic table.

• Nonmetal: an element that is typically a nonconductor; nonmetals are in the top right part of the periodic table.

• Metalloid: an element that has properties of both metals and nonmetals.

Important Groups of Elements

• Representative Elements: the elements in the A groups (1,2, 13-18).

• Transition Metals: the elements in B groups (3-12).• Inner Transition Metals: the two rows of metals

(lanthanides and actinides) set at the bottom of the periodic table.


• Alkali Metals: soft, reactive metals in group 1A.• Alkaline Earth Metals: elements in group 2A. • Halogens (salt formers): reactive nonmetals in group

7A.• Noble Gases: the stable, largely inert, gases in group

8A.


• A molecule is a combination of atoms joined so strongly that they behave as a single particle.

• The simplest molecules are diatomic - they contain two atoms.

Molecules

• If all the atoms in a molecule are the same, the substance is an element.

Elements

• If two or more elements form a molecule, it is a molecular compound.

Molecules

• A molecular formula gives the number of every type of atom in the molecule.• The elements present in the molecule are

identified by their symbols.• A subscript number follows each symbol, giving

the number of atoms of that element present in the molecule; the subscript is omitted if only one atom of the element is present.

• A structural formula shows how the atoms are connected in the molecule.

Molecular Formulas

Molecular Formulas

Molecular Mass

• The relative mass of a molecule in atomic mass units is called the molecular mass of the molecule.

• Because molecules are made up of atoms, the molecular mass of a molecule is obtained by adding together the atomic mass of all the atoms in the molecule.

• The formula for a molecule of water is H2O. This means one molecule of water contains two atoms of hydrogen, H, and one atom of oxygen, O. The molecular mass of water is then the sum of two atomic masses of H and one atomic mass of O:

• MM = 2(at. wt. H) + 1(at. wt. O) • MM = 2(1.01 u) + 1(16.00 u) = 18.02 u


Molecular Mass

• The clear liquid is carbon disulfide, CS2. It is composed of carbon (left) and sulfur (right). What is the molecular weight for carbon disulfide?

• Answer: MW = 1(atomic weight C) + 2(atomic weight S) 12.01 u + 2(32.07 u) = 76.15 u


• One substance present in smog is dinitrogen tetroxide (N2O4). Calculate its molecular mass.

• What is the molecular mass of the fuel propane (C3H8 )?

Example: Calculate Molecular Mass Values

Chemical Bonding

1. Ionic bond: Attractive force that holds ions of opposite charge together.• Involves transfer of e- from one component to the other.• Occurs between positively-charged metal (loses 1 or more e-) and

non-metal atom or molecule (gains 1 or more e-).• Usually satisfies octet rule• Common to inorganic chemistry

2. Covalent bond: Formed by sharing of electrons.• Occurs between:

• Two non-metals• Nonmetal and metalloid• Two metalloids

• Usually satisfies octet rule• Common to organic chemistry

• An ionic compound is composed of cations and anions joined to form a neutral species.

• Each cation is surrounded by several anions and vice versa.

Ionic Compounds

• The formula of an ionic compound is an empirical formula that uses the smallest whole number subscripts to express the relative numbers of ions.

• The relative numbers of ions in the empirical formula balances the charges to zero.• The formula of sodium chloride is NaCl, because the 1+

ions have to be present in a 1:1 ratio.

• The formula of sodium oxide is Na2O, because the charge of the Na+ and O2- ions balance to zero in a 2:1 ratio.

Formulas of Ionic Compounds

• The position of an element in the periodic table can be used to determine the charges of some ions.• The metallic elements in Groups 1A, 2A, 3B, and Al

(Group 3A) all form cations with a charge equal to the Group number.

• The nonmetals in Groups 6A, 7A, and N in group 5A form anions with a charge of 2-, 1- and 3-, respectively.

Formulas of Ionic Compounds

1A 2A 3B 3A 5A 6A 7A

Li+ Be2+ N3- O2- F-

Na+ Mg2+ Al3+ S2- Cl-

K+ Ca2+ Sc3+ Se2- Br-

Rb+ Sr2+ Y3+ I-

Cs+ Ba2+ La3+

Charges on Common Ions

• Write the empirical formulas of the compound formed by(a) the cation of Ca and the anion of Br.(b) the cation of Al and the anion of O.

Example: Ionic Compounds Formulas

• Polyatomic ion: a group of atoms with a net charge that behaves as a single particle.

• Common molecular/polyatomic ions you should memorize:

• OH- (hydroxide), NH4+ (ammonium), SO4

2- (sulfate), SO32-

(sulfite), PO43- (phosphate), NO2

- (nitrite), NO3- (nitrate),

CO32- (carbonate)

• Ionic compounds formed by these ions will have neutral charges.

Polyatomic Ions

Pb(OH)2

NaOH

(NH4+)2SO4

NH3

MgCl2

H3PO4

HBr

HCl

AgNO3

Name Formula Name Formula

Acetate CH3CO2- Nitrate NO3

-

Carbonate CO32- Nitrite NO2

-

Bicarbonate HCO3- Permanganate MnO4

-

Chlorate ClO3- Phosphate PO4

3-

Perchlorate ClO4- Hydrogen

phosphate HPO4

2-

Chromate CrO42- Dihydrogen

phosphate H2PO4

-

Cyanide CN- Sulfate SO42-

Dichromate Cr2O72- Bisulfate HSO4

-

Hydroxide OH- Sulfite SO32-

Common Polyatomic Anions

• Write the formulas of the compounds that contain:(a) the calcium ion and nitrate ion.(b) the ammonium ion and the dichromate ion.

Example: Polyatomic Ions Formulas

• Formula mass is the sum of the atomic masses of all atoms in the empirical formula of an ionic compound.

The formula mass of Ca(NO2)2 is:1(Ca) x 40.08 = 40.082(N) x 14.01 = 28.024(O) x 16.00 = 64.00

Formula mass =132.10 u

Formula Mass of Ionic Compounds

• Chemical nomenclature is the organized system for naming compounds.

• Some of the basic rules of nomenclature are given here for:• Ionic compounds• Acids• Molecular compounds• Organic compounds

Chemical Nomenclature

Naming Ionic Binary Compounds

potassium (K+) + chlorine (Cl-)

[Name of Metal] + [nonmetal stem + ide] =

potassium chloride (KCl)

strontium (Sr2+) + oxygen (O2-) strontium oxide (SrO)

3 calcium (Ca2+) + 2 nitrogen (N3-) calcium nitride (Ca3N2)

Some metals may form more than 1 type of charged ion. Exs: Cu+ and Cu2+; Fe2+ and Fe3+

Compounds with these ions would be named by adding a roman numeral equivalent to charge in parentheses after metal name:

copper (Cu+) + chlorine (Cl-)

[Name of Metal] + [nonmetal stem + ide] =

copper(I) chloride (CuCl)

iron (Fe2+) + 2 chlorine (Cl-) iron(II) chloride (FeCl2)

iron (Fe3+) + 3 chlorine (Cl-) iron(III) chloride (FeCl3)

Anion Name Anion Name

H- Hydride F- Fluoride

N3- Nitride Cl- Chloride

O2- Oxide Br- Bromide

S2- Sulfide I- Iodide

Common Monatomic Anions

Naming ionic compounds containing polyatomic ions

1. Give the name of the metal first.

2. Make sure that charges add up to zero.

3. Put parentheses around polyatomic ions if more than 1 used.

potassium phosphateK and PO43-

Mg and PO43-

K3PO4

sodium nitrateNa and NO3- NaNO3

Mg3(PO43-)2 magnesium phosphate

NO3- and NH4

+ NH4NO3 Ammonium nitrate

Write formulas for the following and name them:

• Name the following ionic compounds:(a) NH4Br (b) Ca(NO3)2 (c) MnSO4

• Give the formula of the following ionic compounds:(a) chromium(III) nitrate(b) potassium sulfate(c) ammonium dichromate

Example: Ionic Compounds Names

• An acid is a compound that produces hydrogen ions when dissolved in water, and for the present can be considered as hydrogen cations combined with one of the anions already discussed.

• For example HCl, HNO3 and H2SO4 are all acids in water solution.

Acids

• If the anion name ends in “ide”, change the ending to “ic” and add the prefix “hydro”. This is followed by the word acid.

Acid Name Anion Name

HBr hydrobromic acid Br- Bromide

H2S hydrosulfuric acid S2- Sulfide

HCN hydrocyanic acid CN- Cyanide

Naming Acids

• If the polyatomic anion name ends in “ate”, change the ending to “ic”; if it ends in “ite” change the ending to “ous”. This is followed by the word acid.

Acid Name Anion Name

H3PO4 Phosphoric acid PO43- Phosphate

HClO4 Perchloric acid ClO4- Perchlorate

HNO2 Nitrous acid NO2- Nitrite

Naming Acids (cont’d)

• Many molecular compounds have nonsystematic common names; e.g. water (H2O), ammonia (NH3), and methane (CH4).

• The systematic names of binary molecular compounds are similar to those of ionic compounds - the name of the first element followed by the name of the second element with the ending “ide”

Names of Molecular Compounds

• The order of the elements in the names and formulas of molecular compounds is:• The element farther to the left in the periodic table appears

first (lesser EN)• The element closer to the bottom within any group is first

(lesser EN).• Hydrogen is first when combined with 6A and 7A elements; it

is named second when combined with groups 1A through 5A elements.

• Oxygen is second, except when combined with fluorine.

Order of Element Names

• Often the same elements form more than one compound. Numerical prefixes are used to give the number of atoms present in the molecule.

Numerical Prefixes in Names

• What are the names of the following compounds?(a) H2SO4

(b) SF6

(c) C3O2

(d) TiO2

Example: Naming Compounds

CO

CO2

H2O

N2O5carbon monoxide

carbon dioxide

dihydrogen monoxide

dinitrogen pentoxide

CCl4 carbon tetrachloride

S2O7 disulfur heptoxide

• Hydrocarbons are organic compounds that contain only the elements hydrogen and carbon. • Alkanes are hydrocarbons that have the general

formula CnH2n+2 (n = integer),

• Cycloalkanes are hydrocarbons that contain a ring of carbon atoms and have the formula CnH2n.

• More complex organic compounds contain functional groups; atoms or small groups of atoms that undergo characteristic reactions. • Alcohol contains the –OH functional group.• Ethers contain the C-O-C functional group.

Organic Compounds

• Alkanes after CH4 (methane) and C2H6 (ethane) are named by using the suffix -ane with the appropriate prefix (pro- for three, but- for four; after that, numerical prefixes are used – pent- for five, hex- for six, etc).

• C5H12 is pentane; C8H14 is octane.

• Cycloalkanes are named the same as with alkanes adding the prefix cyclo.

Alkane Nomenclature

• Alkanes and cycloalkcanes can have substituents. • An alkyl group is an alkane with one hydrogen

removed at the location it is attached as a substituent. Alkyl substituents have the base alkane name with a –yl ending. A methyl group is

–CH3

• Halides like –Cl (chloride) are also substituents.• Substituents are numbered on the chain.

Alkane Nomenclature

• Name the following compounds.

(a) CH3CH2CH2CH3

(b) CH3CHBrCH2CH2CH3

Alkane Nomenclature

• Ionic compounds are usually combinations of metals and nonmetals, while molecular compounds usually contain only nonmetals.

• Ionic compounds are usually hard, brittle solids with high melting points; molecular compounds have lower melting points, and may be liquids or gases at room temperature.

Ionic and Molecular Compounds

• Most ionic compounds dissociate into individual cations and anions when dissolved in water.• NaCl dissociates into Na+ and Cl- in water.

Dissociation of Ionic Compounds

Na Cl+ ClNa

1+ 1-

EN: 0.9 EN: 3.0

0.79 Å 0.91 Å

2.23 Å 0.97 Å

Na Cl

H C

Na Cl

H C

• An electrolyte is a substance that produces ions in water solution.

• Ionic compounds are electrolytes - they conduct electricity when dissolved in water. • Ionic compounds heated until they melt to form a

liquid also conduct electricity.

Electrolytes

• Water and compounds that dissolve in water as neutral molecules are nonelectrolytes, they do not conduct electrical current.

• Salts and other electrolytes in the water will conduct electricity.

• Most molecular compounds (covalently-bonded) are also nonconducting.

Nonelectrolytes

Measured conductivity of (a) ionic solids, (b) melted or (c) dissolved ionic compounds and (d) molecular compounds. Melted or dissolved ionic compounds conduct.

Electrical Conductivity

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Lecture 02 (Chapter 2) Atoms, Molecules, and Ions