Atoms, Molecules, and Ions 2.1 Atomic Theory 2.2 The Structure of the Atom 2.3 Nuclear Structure; Isotopes 2.4 Atomic Weights 2.5 Periodic Table 2.6 Chemical

Atoms, Molecules, and Ions

2.1 Atomic Theory

2.2 The Structure of the Atom

2.3 Nuclear Structure; Isotopes

2.4 Atomic Weights

2.5 Periodic Table

2.6 Chemical Formulas

Copyright © Houghton Mifflin Company.All rights reserved. Presentation of Lecture Outlines, 2–2

Atomic Theory of Matter

• Dalton’s Atomic Theory– All matter is composed of

indivisible atoms. An atom is an extremely small particle of matter that retains its identity during chemical reactions.

– An element is a type of matter composed of only one kind of atom, each atom of a given element having the same properties. Mass is one such property. Thus the atoms of a given element have a characteristic mass.

http://micro.magnet.fsu.edu/primer/java/scienceopticsu/powersof10/


1. A compound is a type of matter composed of atoms of two or more elements chemically combined in fixed proportions.

2. The relative numbers of any two kinds of atoms in a compound occur in simple ratios.

3. Water, for example, consists of hydrogen and oxygen in a 2 to 1 ratio.

4. A chemical reaction consists of the rearrangements of the atoms present in the reacting substances to give new chemical combinations present in the substances formed by the reaction.

5. Atoms are not created, destroyed, or broken into smaller particles by any chemical reaction.


Postulates of Dalton’s Atomic Theory



Particle Structure of the AtomElectron – Discovered in 1898 by J.J. Thompson

Electron : A very light negatively charged subatomic particle. -

+e-

Another view of the Thompson Experiment



The nuclear model of the atom.

Ernest Rutherford, a British physicist, put forth the idea of the nuclear model of the atom in 1911, based on experiments done in his laboratory by Hans Geiger and Ernest Morrison.

http://micro.magnet.fsu.edu/electromag/java/rutherford/





– These experiments showed that the atom consists of two kinds of particles: a nucleus, the atom’s central core, which is positively charged and contains most of the atom’s mass, and one or more electrons.


• The structure of the atom

– Electrons are very light, negatively charged particles that exist in the region around the atom’s positively charged nucleus.


In 1909, U.S. physicist, Robert Millikan had obtained the charge on the electron. The electron’s mass was calculated to be 9.109 x 10-31 kg, which is more than 1800 times smaller than the mass of the lightest atom (hydrogen).




Isotopes are atoms whose nuclei have the same atomic number but different mass numbers; that is, the nuclei have the same number of protons but different numbers of neutrons.

Hydrogen Deuterium Tritium

P P PNN

N



How many protons, neutrons and electrons in each of the following:

protons neutrons electrons23Na14N38Ar35Cl36Cl-156Fe

11 12 7

1177

18 20 18

17 18 1717 19 18

26 30 26

Protons Neutrons Electrons

6 6 6

6 7 6

6 8 6


The fractional abundance is the fraction of a sample of atoms that is composed of a particular isotope. Another name is the weighted average of the atomic mass.



Figure 2.11: Diagram of a Simple Mass Spectrometer, Showing the Separation of Neon

Isotopes


Atomic Weights

Calculate the atomic weight of boron, B, from the following data:

ISOTOPE ISOTOPIC MASS (amu) FRACTIONAL ABUNDANCEB-10 10.013 0.1978B-11 11.009 0.8022

B-10: 10.013 x 0.1978 = 1.9805

B-11: 11.009 x 0.8022 = 8.8314

10.8119 = 10.812 amu

( = atomic wt.)


– Since Dalton could not weigh individual atoms, he devised experiments to measure their masses relative to the hydrogen atom.

Atomic Weights

• Dalton’s Relative Atomic Masses

– Hydrogen was chosen as it was believed to be the lightest element. Daltons assigned hydrogen a mass of 1.

– For example, he found that carbon weighed 12 times more than hydrogen. He therefore assigned carbon a mass of 12.


– Dalton’s atomic weight scale was eventually replaced in 1961, by the present carbon–12 mass scale.

Atomic Weights

• Dalton’s Relative Atomic Masses

– One atomic mass unit (amu) is, therefore, a mass unit equal to exactly 1/12 the mass of a carbon–12 atom.

– On this modern scale, the atomic weight of an element is the average atomic mass for the naturally occurring element, expressed in atomic mass units.


The Periodic Table

In 1869, Dmitri Mendeleev discovered that if the known elements were arranged in order of atomic number, they could be placed in horizontal rows such that the elements in the vertical columns had similar properties.

A tabular arrangement of elements in rows and columns, highlighting the regular repetition of properties of the elements, is called a periodic table.


Figure 2.15: A modern form of the periodic table.


– A period consists of the elements in one horizontal role of the periodic table.

The Periodic Table

• Periods and Groups

– A group consists of the elements in any one column of the periodic table.

– The groups are usually numbered.– The eight “A” groups are called main group (or

representative) elements.


– The “B” groups are called transition elements.

The Periodic Table

• Periods and Groups

– The two rows of elements at the bottom of the table are called inner transition elements.

– Elements in any one group have similar properties.


The Periodic Table

Periods and Groups


– A metal is a substance or mixture that has a characteristic luster and is generally a good conductor of heat and electricity.

The Periodic Table

• Metals, Nonmetals, and Metalloids

– A nonmetal is an element that does not exhibit the characteristics of the metal.

– A metalloid, or semi-metal, is an element having both metallic and nonmetallic properties.


The Periodic Table


Chemical Formulas; Molecular and Ionic Substances

• The chemical formula of a substance is a notation using atomic symbols with subscripts to convey the relative proportions of atoms of the different elements in a substance.

– Consider the formula of aluminum oxide, Al2O3. This formula implies that the compound is composed of aluminum atoms and oxygen atoms in the ratio 2:3.


– A molecule is a definite group of atoms that are chemically bonded together – that is, tightly connected by attractive forces.


• Molecular substances

– A molecular substance is a substance that is composed of molecules, all of which are alike.

– A molecular formula gives the exact number of atoms of elements in a molecule.

– Structural formulas show how the atoms are bonded to one another in a molecule.


Figure 2.18: Molecular and structural formulas and molecular models.

Ionic Bonding-Being Like the Noble Gases

All atoms want to have the same number of electrons as the Noble Gases. The Noble Gases have very stable electron configurations. In order to achieve the same electron configuration as the Noble Gases metal atoms will give up electrons to form positive ions (cations) and non-metal atoms will receive or take additional electrons to become negative ions (anions). IONS are charged particles.

N becomes N-3

Al becomes Al+3

Cl becomes Cl- O becomes O-2

Mg becomes Mg+2Na becomes Na+

The positive and negative ions are attracted to each other electrostatically.

Putting Ions Together

Na+ + Cl- = NaCl

Ca+2 + O-2= CaO Na+ + O-2 = Na2O

Al+3 + S-2 = Al2S

3Ca+2 + N-3 = Ca

3N

2

Ca+2 + Cl- = CaCl2

You try these!

Mg+2 + F- =

NH4

+ + PO4

-3 =

K+ + Cl- =

Al+3 + I- =

Sr+2 + P-3 =

Li+ + Br- =

Sr3P

2

AlI3

MgF2

(NH4)

3PO

4

KCl

LiBr

Not NH43

PO4


Figure 2.21: A model of a portion of crystal.


Atoms, Molecules, and Ions

2.7 Organic Compounds

2.8 Naming Simple Compounds

2.9 Writing Chemical Equations

2.10 Balancing Chemical Equations


– Chemical compounds are classified as organic or inorganic.

Chemical Substances; Formulas and Names

• Naming simple compounds

– Organic compounds are

compounds that contain carbon combined with other elements, such as hydrogen, oxygen, and nitrogen.

– Inorganic compounds are compounds composed of elements other than carbon.


– An important class of molecular substances that contain carbon is the organic compounds.


• Organic compounds

– Organic compounds make up – the majority of all known compounds. – The simplest organic compounds are hydrocarbons, or

compounds containing only hydrogen and carbon.

– Common examples include methane, CH4, ethane, C2H6, and propane, C3H8.



– Most ionic compounds contain metal and nonmetal atoms; for example, NaCl.


• Ionic compounds

– You name an ionic compound

by giving the name of the cation followed by the name of the anion.

– A monatomic ion is an ion formed from a single atom.– Table 2.4 lists some common monatomic ions of the

main group elements.


– Most of the main group metals form cations with the charge equal to their group number.


Rules for predicting charges on monatomic ions

– The charge on a monatomic anion for a nonmetal equals the group number minus 8.


5.7 NomenclatureNaming of Compounds

Binary Compounds have two types of atoms (not diatomic which has only two atoms).

Metals (Groups I, II, and III) and Non-Metals

Metal _________ + Non-Metal _________ideSodium Chlorine

Sodium Chloride NaCl

Metals (Transition Metals) and Non-Metals

Metal ______ +Roman Numeral (__) + Non-Metal ________ide Iron III Bromine

Iron (III) Bromide FeBr3

Compare with Iron (II) Bromide FeBr2


NomenclatureNaming of Compounds

Binary Compounds have two types of atoms (not diatomic which has only two atoms).

Metals (Transition Metals) and Non-MetalsOlder System

Metal (Latin) _______ + ous or ic + Non-Metal ________ide Ferrous Bromine

Ferrous Bromide FeBr2

Compare with Ferric Bromide FeBr3

Non-Metals and Non-MetalsUse Prefixes such as mono, di, tri, tetra, penta, hexa, hepta, etc.

CO2 Carbon dioxide CO Carbon monoxide

PCl3 Phosphorus trichloride CCl4 Carbon tetrachloride

N2O5 Dinitrogen pentoxide CS2 Carbon disulfide


Let’s Practice!Name the following

CaF2

K2S

CoI2

SnF2

SnF4

OF2

CuI2

CuI

SO2

SrS

LiBr

Strontium SulfideLithium Bromide

Copper (I) Iodide or Cuprous Iodide

Sulfur dioxide

Copper (II) Iodide or Cupric Iodide

Oxygen diflourideTin (IV) Flouride or Stannic Flouride

Tin (II) Flouride or Stannous Flouride

Cobalt (II) Iodide or Cobaltous IodidePotassium Sulfide

Calcium Flouride


– A polyatomic ion is an ion consisting of two or more atoms chemically bonded together and carrying a net electric charge.

– Table 2.6 lists some common polyatomic ions. Here a few examples.


• Polyatomic ions

nitrite NO

nitrate NO

2

3

sulfite SO

sulfate SO2

3

24


Ions You Should Know(See Table 2.6 on page 66)

• NH4+ - Ammonium

• OH- - Hydroxide• CN- - Cyanide

• SO42- - Sulfate

• SO32- - Sulfite

• ClO- - Hypochlorite

• ClO2- - Chlorite

• ClO3- - Chlorate

• ClO4- - Perchlorate

• MnO4- - Permanganate

• NO3- - Nitrate

• NO2- - Nitrite

• O22- - Peroxide

• PO43- - Phosphate

• PO33- - Phosphite

• CO32- - Carbonate

• HCO3- - Bicarbonate or

Hydrogen Carbonate

• CrO4-2 - Chromate

• Cr2O7-2 - Dichromate


Acids (with H in front)

Binary acids (without oxygen in formula)

Hydro _________ ic Acid

HCl Hydrochloric acid HBr Hydrobromic acid

Oxy acids (with oxygen in formula)

-ate goes to –ic and –ite goes to -ous

HNO3 Nitric acid HNO2 Nitrous acid

H2SO4 Sulfuric acid H2SO3 Sulfurous acid

H3PO4 Phosphoric acid H3PO3 Phosphorous acid


Lets Practice!

HFNa2CO3

H2CO3

KMnO4

HClO4

H2S

NaOH

CuSO4

PbCrO4

H2O

NH3

Hydrooxic acid (no……just water)

Nitrogen trihydride (no..just ammonia)

Copper (II) sulfate or Cupric sulfate

Lead (II) chromate or Plubous chromate

Sodium hydroxide

Hyrdosulfuric acidPerchloric acid

Potassium permanganate

Sodium carbonate

Hydroflouric acid

Carbonic acid


More Practice

Na2SO4 Na2SO3

Sodium Sulfate Sodium Sulfite

AgCN Cd(OH)2

Silver Cyanide Cadmium Hydroxide

Ca(OCl)2 KClO4

Calcium Hypochlorite Potassium Perchlorate


– A hydrate is a compound that contains water molecules weakly bound in its crystals.


• Hydrates

– Hydrates are named from

the anhydrous (dry) compound,

followed by the word “hydrate”

with a prefix to indicate the number of water molecules per formula unit of the compound.

– For example, CuSO4. 5H2O is known as

copper(II)sulfate pentahydrate.


– The reactants are starting substances in a chemical reaction. The arrow means “yields.” The formulas on the right side of the arrow represent the products.

– A chemical equation is the symbolic representation of a chemical reaction in terms of chemical formulas.

Chemical Reactions: Equations

• Writing chemical equations

NaCl2ClNa2 2

– For example, the burning of sodium and chlorine to produce sodium chloride is written


– In many cases, it is useful to indicate the states of the substances in the equation.

– When you use these labels, the previous equation becomes


• Writing chemical equations

)s(NaCl2)g(Cl)s(Na2 2


Taking 20 kids to the zoo?

What if you came home with only 18 kids?Parents are funny

that way!

What if you came home with 22 kids?

At who’s house would you drop them off?


Balancing using the underline method.

Na2O(s) + H2O(l) NaOH(aq)2

CH4(g) + O2(g) CO2(g) + H2O(g)22

Fe(s) + O2(g) Fe2O3(s)2 24 3

LiOH(s) + CO2(g) LiHCO3(s)

KClO3(s) KCl(s) + O2(g)

MnO2

32 2



• Balance the following equations.

332 POCl PCl O

26424 N OP ON P

232232 SO O As O SAs

24243243 )POCa(H POH )(POCa

2 2

6 6

62 29

34

Documents

Atoms, Molecules, and Ions 2.1 Atomic Theory 2.2 The Structure of the Atom 2.3 Nuclear Structure; Isotopes 2.4 Atomic Weights 2.5 Periodic Table 2.6 Chemical