Chapter 2

Atoms and

Elements

AP CHEMISTRY

Copyright 2011 Pearson Education, Inc.

Dalton’s Atomic Theory

2

1. Each element is composed of tiny, indestructible particles called atoms

2. All atoms of an element have the same mass and properties that distinguish them from atoms of other elements

3. Atoms combine in simple, whole-number ratios to form molecules

4. In a chemical reaction, atoms of one element cannot change into atoms of another element, they simply rearrange the way they are attached to each other

Decide which statement is correct according

to Dalton’s model of the atom

3 1 2 3 4

0% 0%0%0%

1. Copper atoms can combine with

zinc atoms to make gold atoms

2. Bulk water is composed of many

identical molecules, each having

one oxygen atom and two hydrogen

atoms

3. Because the mass ratio of Fe:O in

wüsite is 1.5 times larger than the

Fe:O ratio in hematite, there must

be 1.5 Fe atoms in a unit of wüsite

and 1 Fe atom in a unit of hematite

4. Some carbon atoms weigh more

than other carbon atoms

incorrect; atoms must combine in small whole-number ratios. You

can get the Fe:Fe mass ratio to be 1.5 if the formula for wüsite is

FeO and the formula for hematite is Fe2O3

4

incorrect; atoms of one element cannot turn into atoms of another

element by a chemical reaction

correct; atoms combine together in compounds in small whole

number ratios, so that you could describe a compound by

describing the number of atoms of each element in a molecule.

4. Because the mass ratio of Fe:O in wüsite is 1.5 times larger than

the Fe:O ratio in hematite, there must be 1.5 Fe atoms in a unit of

wüsite and 1 Fe atom in a unit of hematite

The Statements According to Dalton’s Theory are:

1. Copper atoms can combine with zinc atoms to make gold atoms

2. Bulk water is composed of many identical molecules, each having

one oxygen atom and two hydrogen atoms

3. Some carbon atoms weigh more than other carbon atoms

incorrect; all atoms of the same element weigh the same

5

Charge

• Two kinds of charge exist in nature, + and –

• Two opposite charges attract, + attracts –

• Two like charges repel, + repels +, – repels –

• To be neutral, either no charge is present, or equal amounts of + and – charges are present

Cathode Ray Tube (CRT)

6

• An evacuated glass tube with metal electrodes

connected to a high voltage power supply

• A glowing ray emanates from the cathode and

terminates at the anode

What is the ray of a cathode ray tube?

light or charged particles

• Thomson believed that the cathode ray was

composed of tiny electrically charged particles

He designed an experiment to demonstrate this by

measuring the amount of force it took to deflect the

particles path a given amount

A force was exerted by placing a charged electric

field around the CRT

7

1. charged matter is attracted to an electric field

2. light’s path is not deflected by an electric field

8

Thomson’s Experiment

+++++++++++

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

Power Supply- +

Cathode Anode

(-) (+)

Thomson’s Results

• the beam was deflected toward the + plate,

Cathode rays are made of tiny negatively charged

particles, not light

• Every material tested had these same particles

• The amount of deflection is related to two factors,

the charge and mass of the particles

The charge to mass ratio of these particles is

−1.76 x 108 C/g (high charge/mass ratio)

Compared to the charge to mass of the hydrogen

ion, which is

+9.58 x 104 C/g (lower charge/mass ratio)

9

Thomson’s Conclusions

• Since atoms are neutral as elements, their +

and – charges have to cancel

• If the – particle has the same amount of

charge as a + hydrogen ion, then the –

particle must have a mass almost 2000 times

smaller than a hydrogen ion!

• The only way for this to be true is if these

particles were pieces of atoms

apparently, the atom is not unbreakable

10

Thomson’s Conclusions

• Thomson believed that these particles were

therefore the ultimate building blocks of matter

“We have in the cathode rays matter in a new state,

a state in which the subdivision of matter is carried

very much further . . . a state in which all matter . . .

is of one and the same kind; this matter being the

substance from which all the chemical elements are

built up.”

• The cathode ray particles became known as

electrons

tiny, negatively charged particles found in all atoms

11

12

Millikan’s Oil Drop Experiment

(X-rays)

pinhole

13

Millikan’s oil droplet experiment

• An oil droplet that falls into the electric field becomes

negatively charged by the stream of ionizing radiation

When the charge on the drop (in coulombs) / oil drop mass (g)

= acceleration due to gravity / the applied electric field, the

droplet stops falling

q/m = g/E

• The oil drop mass = oil density x oil drop volume This was used to calculate the C/g ratio.

a whole # of electrons was assumed

every drop had a C/g ratio that reduced to a charge of

-1.60 x 1019 C, the charge of the electron

the electron has a mass of 9.1 x 10-28 g

electrons are particles found in all atoms

A New Theory of the Atom

• Because of these experiments, the atom was no

longer believed to be indivisible

• Thomson proposed a new atomic model that

replaced Dalton’s premise that each element is

composed of tiny, indestructible particles (atoms)

The rest of Dalton’s theory was still valid at this point

• Thomson proposed that instead of being a hard,

marble-like unbreakable sphere, the way Dalton

described it, the atom actually had an inner

structure

16

17

J.J. Thomson’s Plum Pudding Model

• Atoms contain many negatively charged electrons held in the atom by their attraction for the atom’s positively charged electric field

The mass of the atom is due to the mass of the electrons

Thomson assumed there were no positively charged pieces in the atom, because none showed up in the cathode ray experiment

the negatively charged particles

should not be near each other because

they would repel

• Atoms are mostly empty space

18

Discovery of Radioactivity

• Henri Becquerel and Marie Curie discovered that certain elements constantly emit small, energetic particles and rays that could penetrate matter

Marie Curie

1867-1934

• Ernest Rutherford discovered that there were three different kinds of emissions

alpha (a) particles with a mass 4x that of the H atom and a positive (+) charge

beta (b) particles with a mass ~1/2000x that of the H atom and a negative (–) charge

gamma (g) rays that are energy rays, not particles

Rutherford’s Experiment to determine

if an atom is mostly empty space

Devised a method to shoot alpha particles through

atoms to see the result

• a particles have a mass of 4 amu & a charge of +2 c.u.

large compared to electrons, so electrons will not deflect

them

• used large target atoms of gold, mass = 197 amu

used very thin sheets of gold so the “bullet” would not be

absorbed

19

20

Rutherford’s Experimental Results

and Conclusions

• > 98% of the a particles went straight through

Atoms are mostly empty space because >98% of the

particles went straight through

• ~ 2% of the a particles went through the foil but

were deflected by large angles

The dense particle is positively charged which

explains the large deflections of ~2% of the particles

• ~ 0.005% of the a particles bounced off the foil

Atom contain a dense particle that is small in volume

compared to the atom but large in mass because ~

0.005% of the particles bounced back

21

22

.

.

.

Nuclear Atom

•

•

••

• ••

• ••

•

•

•

•

• •

••

•

••

•

Plum Pudding

Atom

If atom was like

a plum pudding,

all the a particles

should go

straight through

Almost all a particles

go straight through

Some a particles

go through, but are deflected due to

+:+ repulsion from the nucleus

A few of the

a particles

do not go through

23

Rutherford’s Interpretation

The Nuclear Model

1. Atoms contain a tiny dense center, the nucleus

2. The nucleus contains the atom’s entire mass the electron’s comparative weight/mass is negligible

3. The nucleus is positively charged and it balances the negative charges of the electrons

4. The electrons are dispersed in the empty space of the atom surrounding the nucleus

Drop a pea into any large sports stadium. The pea is to the

nucleus as the stadium is to the atom.

the scale of comparison is called a relative scale

Structure of the Nucleus

• Rutherford proposed that the nucleus had a particle that had the same amount of charge as an electron, but opposite in sign, called protons

• Protons are found in the nucleus and have a charge=+1.60 x1019 C and mass=1.67262 x 10−24 g

Each proton has a charge of +1 charge units (cu)

protons and electrons have equal amounts of charge, but opposite in sign, therefore

atoms must have equal numbers of protons and electrons to be neutral

25

Something is Missing

• How can an atom’s nucleus have multiple protons

with + charges placed in close proximity?

they should repel each other

• Beryllium atoms have four protons

it should weigh 4 amu; but it actually weighs 9.01 amu

the protons account for only 4 amu

Be has four electrons which weigh 4x(~0.00055 amu), far

less than the extra 5 amu needed for mass balance

Where is the extra mass coming from?

26

Proposal to Explain the Difference

• Rutherford and Chadwick proposed that there

was another particle in the nucleus – called a

Neutron

• Neutrons are subatomic particles with a mass

= 1.67493 x 10−24 g

1 amu (slightly heavier than a proton)

• They have no charge

• They are found in the nucleus

27

Atomic Mass Units• We take 1/12th the mass of the carbon atom (6 protons

and 6 neutrons), and call it 1 atomic mass unit (amu)

protons and neutrons have a mass of just over 1 amu

electrons have a mass of 0.00055 amu

too small to be relevant to the total mass of the atom

28

29

Why Does Matter Appear Continuous

If the Atom Is Mostly Empty Space?

The emptiness of the atom is on such a

small scale that the variations in density

cannot be seen

Consider a ladder framework that is

mostly empty (right). This framework, if

large enough, will look solid from a

distance. Also, it has some rigidity based

on the interaction of adjacent latices.

30

Elements

• Each element has a unique number of protons

in its nucleus

The number of protons in the nucleus is called the

atomic number, “Z”

The number of protons defines the element

• All the elements are arranged on the Periodic

Table in order of their atomic numbers

Each element has a unique name and symbol

Each symbol is either one or two letters– one capital letter alone, or

– one capital letter followed by one lowercase letter

31

The Periodic Table of the Elements

Some symbols come from the element ‘s name, like C for carbon. Others come from

the Latin name, like gold (Au-aurum) and copper (Cu-cuprium), or the Greek, like

chlorine (Cl-chloros) and argon (Ar-argos).

32

The atomic number tells you

how many protons are in the

nucleus and how many

electrons are in the atom

33

The Nucleus

• Atomic number (Z) = # of protons

“Z” does not change for each element

Does change from element to element

• Mass Number (A) = # protons + # neutrons

“A” does change for elements and their isotopes

• Both Z and A are whole numbers

34

Isotopes

• It was discovered that the same element could have atoms with different masses called isotopes

• All isotopes of an element have the same number

of protons and electrons

• All isotopes of an element are chemically identical

undergo the exact same chemical reactions

• Isotopes of an element have different masses

Due to different numbers of neutrons

Isotopes are identified by their mass numbers, the

sum of all the protons and neutrons in the nucleus

35

Isotopes

• The observed atomic mass of any element is a weighted average of the weights of all the naturally occurring isotopes there are two isotopes of chlorine found in nature, one

that has a mass of about 35 amu and another that weighs about 37 amu

the percentage of each isotope is called the isotope’s natural abundance

the observed atomic mass of chlorine is 35.45 amu

36

Neon

9.25%221210Ne-22 or

0.27%211110Ne-21 or

90.48%201010Ne-20 or

Percent

Natural

Abundance

A, Mass

Number

Number of

Neutrons

Number of

ProtonsSymbol

Ne2010

Ne2110

Ne2210

Which is not an example of isotopes?

37 1 2 3 4

0% 0%0%0%

1. Z=1, N=1 and Z=1, N=2

2. Z=9, N=9 and Z=8, N=9

3. Z=9, N=9 and Z=9, N=10

4. Z=1, N=3 and Z=1, N=2

Example 2.3b: How many protons, electrons,

and neutrons are in an atom of ?

for most stable isotopes, n0 ≥ p+Check:

Z = 24 = # p+ = # e−Solution:

in neutral atom, # p+ = # e-

mass number = # protons + # neutrons

Conceptual

Plan:

Relationships:

therefore A = 52, Z = 24

# protons, # electrons, # neutrons

Given:

Find:Cr52

24

symbol atomic

number# p+ # e−

symbol atomic number

& mass number# n0

A = Z + # n0

52 = 24 + # n0

52-24 = # n0

28 = # n0

38

Cr5224

Complete the table and add all the numbers in the

Electrons column. They add up to ?????

391 2 3 4

0% 0%0%0%

Al2713

1. 122

2. 116

3. 106

4. 98

Complete the table

40

Al2713

C136

Mo9642

Cs13355

116

41

Reacting Atoms

• Elements that undergo chemical reactions do not turn

into other elements

All the atoms present at the start will be there at the end

The number of protons determines the element

the number of protons don’t change in a chemical reaction

• Reactions involve sharing or transfer of electrons

• When atoms gain or lose electrons, they acquire a

charge to become ions

Atoms that gain electrons become negatively charged

anions

Atoms that lose electrons become positively charged

cations

Ions and Compounds

• Ions behave differently than neutral atoms

metallic sodium, made of neutral Na atoms, is

highly reactive and quite unstable

sodium cations, Na+, are found in table salt and

are very nonreactive and stable

• Because materials such as table salt are

neutral overall, there must be a balance

between the charges of cations and anions

42

43

Atomic Structures of Ions

• Nonmetals form anions

• For each negative charge, the ion has one more electron than the neutral atom

Anions are named by changing the ending of the element’s name to -ide

fluorine F + 1e− F− fluoride ion

oxygen O + 2e− O2− oxide ion

F = 9 p+ and 9 e− F− = 9 p+ and 10 e−

P = 15 p+ and 15 e− P3− = 15 p+ and 18 e−

44

Atomic Structures of Ions

• Metals form cations

• For each positive charge, the ion has one less

electron than the neutral atom

• Cations are named the same as the

elemental metal

sodium Na Na+ + 1e− sodium ion

calcium Ca Ca2+ + 2e− calcium ion

Na atom = 11 p+ and 11 e− Na+ ion = 11 p+ and 10 e−

Ca atom = 20 p+ and 20 e− Ca2+ ion = 20 p+ and 18 e−

Complete the table, add the ion charge column, and provide

your answer.

45A. B. C. D. E.

0% 0% 0%0%0%

A. +2

B. -2

C. 0

D. +3

E. -3

3Al

Complete the table

46

2Mg

2S

Br

3Al

47

Mendeleev

• Periodic Law – Saw a repeating pattern of properties when the elements are arranged in order of increasing atomic mass

• Ordered elements in rows by increasing atomic mass from left to right

• Put elements with similar properties in the same column, with heavier mass lower in the column

• Used pattern to predict properties of undiscovered elements

48

Periodic Pattern

49

About ¾ of the elements are classified as metals.

They have a reflective surface, conduct heat and

electricity better than other elements, and are

malleable and ductile

Most of the remaining elements are classified as

nonmetals. Their solids have a non-reflective

surface, do not conduct heat and electricity well,

and are brittle.

A few elements are classified as metalloids.

Their solids have some characteristics of metals

and some of nonmetals.

50

Metals

• All metals are solids at room temperature except Hg which is liquid

• Have reflective, shiny surface

• Conduct heat & electricity

• Are malleable (can be shaped)

• Are ductile (can be drawn or pulled into wires)

• Lose electrons to form cations in reactions

• Comprise ~75% of the elements

• Found on the lower left of the periodic table

51

Nonmetals

• Found in all three states

• Poor conductors of heat &

electricity

• Are brittle

• Gain electrons in reactions to

become anions

• Found in upper right quadrant

of the Periodic table

Hydrogen, a nonmetal) is an

exception

Sulfur, S(s)

Bromine, Br2(l)

Chlorine, Cl2(g)

52

Metalloids

• Show some properties of metals and some of

nonmetals

• Also known as semiconductors

Properties of Silicon

shiny

conducts electricity

does not conduct heat well

brittle

53

The Modern Periodic Table

• Columns are called Groups

Elements with similar chemical and physical

properties are in the same group

designated by a number and letter at top

• Rows are called Periods

Each period shows a pattern of properties

repeated in the period above or below

54

The Modern Periodic Table

• Main group elements have letter “A” designation

• Transition elements have letter “B” designation

All transition elements are metals

• The 2 bottom rows are called inner transition elements

aka rare earth elements

are metals

really belong in Period 6 & 7

55

56

= Alkali metals

= Alkali earth metals

= Noble gases

= Halogens

= Lanthanides

= Actinides

= Transition metals

57

Hydrogen• Nonmetal

• Colorless, diatomic gas

very low melting point, boilng point and density

• Reacts with nonmetals to form molecular

compounds

Reacts with Cl2 to form HCl, an acidic gas

HCl dissolves in water to form acids

Reacts with O2 to form H2O, a liquid

• Reacts with metals to form hydrides

metal hydrides react with water to form H2

• Placed in Group IA

but based on properties, it does not belong there

lithium

sodium

potassium

rubidium

cesium

58

Group IA = Alkali Metals• Low density

• Low melting point

• Soft

• Very reactive

never found in elemental form

• React with water to form basic (alkaline) solutions and H2

2 Na(s) + 2 H2O(l) 2 NaOH(aq) + H2(g)

releases a lot of heat

• Tend to form water-soluble salts

• Salts isolated from seawater are melted and electrolyzed to form elemental form

• Flame tests Li = red, Na = yellow, K = violet

59

Group IIA = Alkali Earth Metals

• Harder, higher melting, and denser metals than

alkali metals

• Reactive, but less reactive than alkali metals

• Form stable, insoluble oxides

Metals are extracted from the oxides

• Oxides are basic (alkaline earth)

• Like IA, IIA metals react with water to form H2

Be = no rxn (NR); Mg = rxn with steam; Ca, Sr, Ba = rxn

with cold water

“rxn” abbreviation for reaction

• Flame tests Ca = red, Sr = red, Ba = green

60

Group VIIA = Halogens

• Nonmetals

Very reactive

React with metals to form ionic

compounds

Cl2, Br2 react slowly with water

Br2 + H2O HBr + HOBr

• All diatomic (X2)

F2 and Cl2 are gases

Br2 liquid

I2 solid

• HX all acids

HF weak < HCl strong< HBr < HI

bromine

iodine

chlorine

fluorine

astatine

61

Group VIIIA = Noble Gases

• All gases at room temperature

very low melting and boiling points

• Very unreactive, practically inert

• Very hard to remove an electron from or

give an electron to a noble gas

62

Ion Charge and the Periodic Table

• The charge on an ion can often be determined from an element’s position in the Periodic Table

• Metals always form positively charged species, cations

For many main group metals, the charge equals the group number

• Nonmetals form negatively charged species, anions

the charge equals the group number

63

potassium cation

sulfide anion

calcium cation

bromide anion

aluminum cation

Group IA K+

Group VIA S2−

Group VIIA Br−

Group IIA Ca2+

Group IIIA Al3+

64

Counting Atoms

• a mole is the mass of a material that contains

6.022 x 1023 atoms/particles/ things

Avogadro’s Number

• the mole’s value is referenced to the number of

atoms in exactly 12 grams of pure carbon-12

1 atom of C-12 weighs exactly 12 amu

the average atomic mass of C is12.01 amu, therefore

1 mole of C atoms weighs 12.01 g

12.01 g contains 6.022 x 1023 atoms

• this provides a relationship between mass and the

number of atoms

65

Counting Atoms by Moles

The number of atoms, 6.022 x 1023,

per mole can be used as a conversion

factor. Given the number of moles of

a substance, you can calculate the

number of atoms, and visa versa

66

Example 2.6: Calculate the number of

atoms in 2.45 mol of copper

because atoms are small, the large number of

atoms makes sense

1 mol = 6.022 x 1023 atoms

2.45 mol Cu

atoms Cu

Check:

Solution:

Conceptual

Plan:

Relationships:

Given:

Find:

mol Cu atoms Cu

A silver ring contains 1.10 x 1022 silver atoms. How many

moles of silver are in the ring?

1 mol = 6.022 x 1023 atoms

67 1 2 3

0% 0%0%

1. 6.62x1045 moles

2. 0.0183 moles

3. 54.7 moles

A silver ring contains 1.1 x 1022 silver atoms. How

many moles of silver are in the ring?

68

because the number of atoms given is less than

Avogadro’s number, the answer makes sense

1 mol = 6.022 x 1023 atoms

1.1 x 1022 atoms Ag

moles Ag

Check:

Solution:

Conceptual

Plan:

Relationships:

Given:

Find:

atoms Ag mol Ag

69

Relationship Between Moles and Mass

• Molar Mass = the mass of 1 mole of atoms

6.022x1023 particles

• The molar mass of an element, in grams, is

numerically equal to the element’s atomic mass,

in amu

the lighter the atom, the less a mole weighs

the more atoms there are in 1 g

• Molar mass can be used as a conversion factor

to convert mass to moles: g x mol/g = mole

• or moles to mass: mol x g/mol = g

70

Mole and Mass Relationships

1 mole

sulfur

32.06 g

1 mole

carbon

12.01 g

71

Example 2.7: Calculate the moles of carbon

in 0.0265 g of pencil lead

because the given amount is much less

than 1 mol C, the number makes sense

1 mol C = 12.01 g

0.0265 g C

mol C

Check:

Solution:

Conceptual

Plan:

Relationships:

Given:

Find:

g C mol C

Calculate the moles of sulfur in 57.8 g of sulfur

1 mol S = 32.07 g

72A. B. C. D.

0% 0%0%0%

A. 1.80

B. 0.555

C. 1850

D. 1.802

Calculate the moles of sulfur in 57.8 g of sulfur

73

because the given amount is much less than 1

mol S, the number makes sense

1 mol S = 32.07 g

57.8 g S

mol S

Check:

Solution:

Conceptual

Plan:

Relationships:

Given:

Find:

g S mol S

74

Example 2.8: How many copper atoms are in

a penny weighing 3.10 g?

because the given amount is much less

than 1 mol Cu, the number makes sense

1 mol Cu = 63.55 g, 1 mol = 6.022 x 1023

3.10 g Cu

atoms Cu

Check:

Solution:

Conceptual

Plan:

Relationships:

Given:

Find:

g Cu mol Cu atoms Cu

How many aluminum atoms are in a

can weighing 16.2 g?

75

because the given amount is much less than 1

mol Al, the number makes sense

1 mol Al = 26.98 g, 1 mol = 6.022 x 1023

16.2 g Al

atoms Al

Check:

Solution:

Conceptual

Plan:

Relationships:

Given:

Find:

g Al mol Al atoms Al

76

Mass Spectrometry

• Mass and isotope abundance are measured with

a mass spectrometer

• Atoms or molecules are ionized, then accelerated

down a tube

some molecules are broken into fragments during

the ionization process

these fragments can be used to help determine the

structure of the molecule

• Their path is bent by a magnetic field, separating

them by mass

77

Mass Spectrometer

78

Mass Spectrum

• A mass spectrum is a chart that gives the relative mass and relative abundance of each particle present

79

Example 2.5: If copper is 69.17% Cu-63 with a mass of 62.9396 amu and the rest Cu-65 with a mass of 64.9278 amu,

find copper’s atomic mass

the average is between the two masses,

closer to the major isotope

Check:

Solution:

Conceptual

Plan:

Relationships:

Cu-63 = 69.17%, 62.9396 amu

Cu-65 = 100-69.17%, 64.9278 amu

atomic mass, amu

Given:

Find:

isotope masses,

isotope fractionsavg. atomic mass

Practice – Ga-69 with mass 68.9256 amu and abundance of

60.11% and Ga-71 with mass 70.9247 amu and abundance of

39.89%. Calculate the atomic mass of gallium.

the average is between the two masses,

closer to the major isotopeCheck:

Solution:

Conceptual

Plan:Relationships:

Ga-69 = 60.11%, 68.9256 amu

Ga-71 = 39.89%, 70.9247 amu

atomic mass, amu

Given:

Find:

isotope masses,

isotope fractionsavg. atomic mass

80

81

82

Scanning Tunneling Microscope

• Gerd Bennig and Heinrich Rohrer found that as you pass a sharp metal tip over a flat metal surface, the amount of current that flows varies with distance between the tip and the surface

• Measuring this “tunneling” current allowed them to scan the surface on an atomic scale – essentially taking pictures of atoms on the surface

83

Operation of a STM

84

Scanning Tunneling Microscope

• Later scientists found that not only can you see the atoms on the surface, but the instrument allows you to move individual atoms across the surface