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General, Organic, and

Biological ChemistryFourth Edition

Karen Timberlake

Chapter 3Atomic Theory and

the Periodic Table

Student notes

© 2013 Pearson Education, Inc.

© 2013 Pearson Education, Inc. Chapter 3, Section 1

Elements are pure substances from which all other

things are built.

gold carbon aluminum

2

The Elements

© 2013 Pearson Education, Inc. Chapter 3, Section 1 3

Sources of Some Element Names

Some elements are named for planets, mythological

figures, minerals, colors, scientists, and places.

© 2013 Pearson Education, Inc. Chapter 3, Section 1 4

A symbol

represents the name of an element.

consists of 1 or 2 letters.

starts with a capital letter, 2nd letter always lower-

case.

Examples:

1-Letter Symbols 2-Letter Symbols

C carbon Co cobalt

N nitrogen Ca calcium

F fluorine Al aluminum

O oxygen Mg magnesium

Symbols of Elements


The Periodic Table

5

• First proposed by Russian chemist Dimitry Mendeleev in 1869, modified

and today looks like this


Periods and Groups

Mendeleev’s table was based on the periodic repetition

of the properties of the elements and listed the

elements in order of atomic weights, today based on

Atomic number

On the periodic table,

groups contain elements with similar properties and

are arranged in vertical columns ordered from left to

right. Also called families

periods are the horizontal rows of elements, and

they are counted from the top as Period 1 to

Period 7.

6


Periods and Groups

7


Groups

Group numbers

numbers to identify the columns from left to right.

the letter A for the representative elements (1A to 8A)

and the letter B for the transition elements.

Newer system uses numbers from 1-15

The representative, or main group, elements

include the first 2 groups, 1A (1) and 2A (2), in

addition to groups 3A (13), 4A (14), 5A (15), 6A (16),

7A (17), and 8A (18).

Some groups have common names : 1A = alkali

metals, 2A = alkaline earth metals, 7A = Halogens,

8A = the noble gases

8


Main Group Elements

9

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10

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3 Categories of Elements -- Metals,

Nonmetals, and Metalloids

A heavy zigzag (stairstep)

line separates the metals

from the nonmetals.

Metals (blue) are located

to the left of the line.

Nonmetals (yellow) are

located to the right.

Metalloids (green) are

located along the heavy

zigzag line between the

metals and nonmetals

(have properties of both).

11


Properties of Metals, Nonmetals,

and Metalloids

Metals are

shiny and ductile.

good conductors of heat and electricity.

Nonmetals are

not especially shiny, ductile, or malleable.

poor conductors of heat and electricity.

Metalloids are

better conductors than nonmetals, but not as good as

metals.

used as semiconductors and insulators.

12



Karen Timberlake

Chapter 3The Atom

© 2013 Pearson Education, Inc.

A Brief History of

Atomic Theory


John Dalton’s Atomic Theory ( circa

1804)

14

Dalton theorized that Atoms

are tiny particles of matter too small to see,

are able to combine with other atoms to make compounds, and

are similar to each other for each element and different from atoms of other elements.

A chemical reaction is the rearrangement of atoms.

Dalton envisioned

atoms to be solid,

indivisible spheres, like

billiard balls called the

“billiard ball model”


Atomic Theory in the late 1890’s

Discovery of radioactivity and the discovery of

the first subatomic particle (the electron)

meant model had to change.

JJ Thomson, discoverer of the electron,

developed “plum pudding model.”

Electron was tiny (1/2000th the size of the

atom), negatively charged particle

As atom electrically neutral, electron must be

embedded in “positive dough” of atom like

plums in plum pudding

15


Rutherford’s Gold-Foil

Experiment (1911)

While exploring the behavior of thin sheets of metals

when bombarded with alpha particles (+ charged

particles emitted by radiactive atoms) Ernest

Rutherford’s gold-foil experiment revealed that when

these + charged particles were aimed at atoms of

gold

most went straight through the atoms, but

Occasionally, some were deflected

Conclusion:

There must be a small, dense, positively charged

core (nucleus) in the atom that deflects positive

particlesthat come close.

16


Rutherford’s Gold-Foil

Experiment

17

(a) Positive particles are aimed at a piece of gold foil. (b) Particles that come

close to the atomic nuclei of gold are deflected from their straight path.


The Nuclear Model of the Atom

The atom is mostly empty space

All of the positive charge is located in a tiny,

dense nucleus

The negative electrons are located at a

distance away and must be constantly

moving to avoid being pulled into the nucleus

18


Discovery of Proton and Neutron

Positive charge comes in nucleus actually

due to a particle, called the proton

(Rutherford, 1919)

More mass in the nucleus than protons could

account for in 1932, an electrically neutral

particle called the “neutron” was discovered

by James Chadwick.

19


The Bohr Model (1913)

Proposed by Danish physicist Niels Bohr

Problems with Rutherford’s model as

conflicted with laws of physics

Bohr proposed new laws were needed for tiny

particles like electrons led to development

of quantum physics

Bohr’s model solved some of these problems

Main ideas electrons can only have certain

allowable energies, which correspond to

different distances from the nucleus = Energy

Levels20



Energy levels radiate away from nucleus

Energy levels are labeled by what is called

the principal quantum number “n”

Each holds a distinct number of electrons

which corresponds to 2n2

n = 1 holds 2(1)2 = 2 electrons



21



22

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Modern Atomic Theory

Based on Quantum Physics, which was

developed in the 1920s

Treats the electron as both a particle and a

standing wave

As in the Bohr model, the electron can have

only certain allowable energies (energies of

e- are quantized) (energy levels)

Solutions to the math equations of quantum

physics provide the most probable region

around the nucleus of finding an electron.

These “probabilibty regions” are also known

as orbitals 23


Structure of the Atom

An atom consists of

a nucleus that

contains protons

and neutrons, and

electrons in a

large, empty space

around the

nucleus.

24


Subatomic Particles

Atoms contain subatomicparticles such as

Protons, which have a positive (+) charge;

electrons, which have a negative (–) charge; and

neutrons, which have no charge.

Experiments show that like charges repel and unlikecharges attract.

25


Atomic Mass Scale

By the 1860’s, chemists had devised a relative mass

scale for atomic weights, or masses, today, this is

called the atomic mass

On the atomic mass scale for subatomic particles,

1 atomic mass unit (amu) is defined as 1/12 of the mass

of the carbon-12 atom. Therefore,

a proton has a mass of about 1 (1.007) amu.

a neutron has a mass of about 1 (1.008) amu.

an electron has a very small mass, 0.00055 amu.

1 amu = 1.66 x 10-24 g

26


Particles in the Atom

27

• We are going to round off the mass of the

proton and the neutron to 1.00 amu each

• Remember, 1 amu = 1.66 x 10-24 g, that’s

why we use amu’s for atomic masses instead

of grams!


Atomic Number

The atomic number

is specific for each element.

is the same for all atoms of an element.

is equal to the number of protons in an atom.

appears above the symbol of an element in the

periodic table.

28

11

Na

Atomic Number

Symbol


Atomic Number and Protons

Each element has a unique atomic number equal to the

number of protons:

Hydrogen has atomic number 1; every H atom has

one proton.

Carbon has atomic number 6; every C atom has six

protons.

Copper has atomic number 29; every Cu atom has

29 protons.

29


Number of Electrons in an Atom

All atoms of an element are electrically neutral; they

have

a net charge of zero.

an equal number of protons and electrons.

Number of protons = Number of electrons

Example:

Aluminum atoms have 13 protons and 13 electrons; the

net charge is zero.

30


Mass Number

The mass number represents the number of subatomic

particles in the nucleus, which is equal to the sum of the

number of protons + number of neutrons.

Since protons and neutrons account for the majority of

mass in an atom, we call this the mass number.

31




Karen Timberlake

3.5

Isotopes and

Atomic Mass

Chapter 3Atoms and Elements

© 2013 Pearson Education, Inc.Lectures


Isotopes Discovery of neutron led to realization that atoms of

the same element are not all identical some have

more neutrons than others

Isotopes

are atoms of the same element that have different

mass numbers.

have the same number of protons but different

numbers of neutrons.

can be distinguished by atomic symbols.

33


23

Isotopes and Mass and Atomic Symbols

(Nuclear, or Isotopic, Notation)

Since each isotope of an element has a different

number of neutrons, each isotope’s mass number will

be different. We write these as atomic symbols:

Mass numbers are in the upper left corner.

Atomic numbers are in the lower left corner.

Example: An atom of sodium with atomic number 11

and a mass number 23 has the following atomic

symbol:

mass number

atomic number

34

11Na


Atomic Symbols

(nuclear/isotopic notation)

For an atom, the atomic symbol gives the number of

protons (p+),

neutrons (n), and

electrons (e–).

8 p+ 15 p+ 30 p+

8 n 16 n 35 n

8 e– 15 e– 30 e–

35

168 O 31

15P 6530

Zn


Isotopes of Magnesium

36


Average Atomic Mass

37

The average atomic mass of an element

is listed below the symbol of each

element on the periodic table.

gives the mass of an “average” atom of

each element compared to C-12.

is not the same as the mass number.

is calculated using a weighted average.

Na

22.99


Some Elements and Their Average

Atomic Masses

38

Most elements have two or more isotopes that

contribute to the atomic mass of that element.


Atomic Mass for Cl

The atomic mass of chlorine is

based on all naturally

occurring Cl isotopes.

not a whole number.

the weighted average

of the Cl-35 and Cl-37

isotopes.

39


Calculating the Atomic Mass for Cl

40

To calculate the atomic mass of an element, we

need to know the percent abundance of each

isotope and its mass.




Karen Timberlake

3.6

Electron Arrangement

in Atoms




Electrons and the Properties of an

Element

Remember, Atoms contain

a very small nucleus packed with neutrons and

positively charged protons, contributes most to the

mass of an atom

a large volume of space around the nucleus that

contains the negatively charged electrons.

Big IdeaIt is the electrons that determine the physical and

chemical properties of atoms.

So, we must learn more about the electronic

structure of the elements.

42


Electron Energy Levels

Electrons surround the nucleus in specific energy

levels

Each energy level has a principal quantum number

(n).

The lowest energy level, which is closest to the

nucleus, is labeled n = 1.

The second-lowest energy level is labeled n = 2, the

third n = 3, and so on.

43


Electron Energy Levels

44

Electron energy levels increase

in energy and number as

electrons get farther away from

the nucleus.

The higher the electron energylevels,

the more electrons they hold.

the more energy the electrons have.


Sublevels

Within each energy level, we have sublevels that

contain electrons with identical energy.

are identified by the letters s, p, d, and f.

According to the mathematics of quantum theory, the

number of sublevels within a given energy level

is equal to the value of the principal quantum number, n.

So, n= 1 has 1 sublevel -- the 1s sublevel

n=2 has 2 sublevels – the 2s and the 2p sublevel

n=3 has 3 sublevels – the 3s, 3p, and 3d sublevel

n=4 has 4 sublevels – the 4s, 4p, 4d, and 4f sublevel

45


Sublevels

n= 1 has 1 sublevel -- the 1s sublevel

n=2 has 2 sublevels – the 2s and the 2p sublevel

n=3 has 3 sublevels – the 3s, 3p, and 3d sublevel

n=4 has 4 sublevels – the 4s, 4p, 4d, and 4f sublevel

Each sublevel designation (s, p, d, f) has a maximum

number of electrons that can be accommodated in

orbitals.

Each orbital can only hold 2 electrons max, due to

repulsions.

Each sublevel has a specific number of orbitals

(regions in space) associated with the location of

electrons. 46


Energy Levels , Sublevels and

Orbital types

47

The s sublevel = 1 orbital shape, represented by the yellow

box below

The p sublevel = 3 orbital shapes (3 green boxes)

The d sublevel = 5 orbital shapes (5 salmon boxes) The f sublevel = 7 orbital shapes (7 lavender boxes)


Energy of Sublevels

Within any energy level,

the s sublevel has the lowest energy.

the p sublevel follows and is slightly higher in energy.

the d sublevel follows the p and is slightly higher in

energy than the p.

the f sublevel follows the d and is slightly higher in

energy than the d.

48


Orbitals

49

Each electron sublevel consists of orbitals, which

are regions where there is the highest probability

of finding an electron.

have their own unique three-dimensional shape.

Each can hold up to 2 electrons.


s Orbitals

50

We know that s orbitals have a spherical shape,

centered around the atom’s nucleus (located at the

origins of the xyz axis shown below.

Only one orientation a sphere can have in 3-d

space, so only one type of s orbital.

The s orbitals get bigger

as the principal quantum

number, n, gets bigger.

The s orbitals can hold

up to 2 electrons

that must spin in

opposite directions.


p Orbitals

51

There are three p orbitals in each energy level,

starting with energy level 2. They

have a two-lobed shape, much like tying a balloon

in the middle, and can hold 2 electrons each.

They are oriented along the axes of the 3-d graph

and are labeled x, y, and z.

increase in size as the value of n increases.


d Orbitals

52

There are five p orbitals in each energy level,

starting with energy level 3. They

Four of them have a 4 leaf clover shape

One looks like a dumbell with a donut around it.

increase in size as the

value of n increases.

Each can hold a max. of

2 e-, spinning in opposite

directions


Sublevels and Orbitals

Each sublevel consists of a specific number of

orbitals.

An s sublevel contains one s orbital.

A p sublevel contains three p orbitals.

A d sublevel contains five d orbitals.

An f sublevel contains seven f orbitals.

53


Electron Capacity in Sublevels

54




Karen Timberlake

3.7

Orbital Diagrams and

Electron Configurations




Order of Filling

Energy levels are filled with electrons

in order of increasing energy.

beginning with quantum number n = 1.

beginning with s followed by p, d, and f in each

energy level.

56


Energy Diagram for Sublevels

57


Orbital Diagrams

58

An orbital diagram shows

orbitals as boxes in each sublevel.

electrons in orbitals as vertical arrows.

electrons in the same orbital with opposite spins (up

and down vertical arrows).

Example:

Orbital diagram for Li

1s2

filled

2s1

half-filled2p

empty


Order of Filling

Electrons in an atom

fill the lowest energy level and orbitals first,

fill orbitals in a particular sublevel with one electron

each until all orbitals are half full, and then

fill each orbital using electrons with opposite spins.

59


Writing Orbital Diagrams

60

The orbital diagram for

carbon has 6 electrons:

2 electrons are used to fill

the 1s orbital.

2 more electrons are

used to fill the 2s orbital.

1 electron is used in two

of the 2p orbitals so they

are half-filled, leaving one

p orbital empty.

Electron

arrangements

in orbitals in

energy levels 1

and 2.


Electron Configuration

An electron configuration

lists the filled and partially filled energy levels in order

of increasing energy.

lists the sublevels filling with electrons in order of

increasing energy.

uses superscripts to show the number of electrons in

each sublevel.

for neon is as follows: number of electrons = 10

1s22s22p6

61


Period 1 Configurations

62

In Period 1, the first two electrons enter the 1s orbital.



63

In Period 2,

lithium has 3 electrons –2 in the 1s and 1 in the 2s.

beryllium has 4 electrons –2 in the 1s and 2 in the

2s.

boron has 5 electrons –2 in the 1s, 2 in the 2s, and

1 in the 2p.

carbon has 6 electrons –2 in the 1s, 2 in the 2s, and

2 in the 2p.


Abbreviated Configurations

In an abbreviated configuration,

the symbol of the noble gas is in brackets,

representing completed sublevels.

the remaining electrons are listed in order of their

sublevels.

Example: Chlorine has the following configuration:

1s22s22p63s23p5

[Ne]

The abbreviated configuration for chlorine is

[Ne]3s23p5.

64



65



66


Electron Configurations and the

Periodic Table

The periodic table consists of sublevel blocks

arranged in order of increasing energy.

Groups 1A and 2A = s block

Groups 3A to 8A = p block

Transition Elements

(This sublevel is (n-1), 1 less

than the period number.) = d block

Lanthanides/Actinides

(This sublevel is (n-2), 2 less

than the period number.) = f block

67


Sublevel Blocks

68


Guide to Using Sublevel Blocks

69


Writing Electron Configurations

Using the periodic table, write the electron configuration

for silicon.

Solution:

Period 1 1s block 1s2

Period 2 2s → 2p blocks 2s2 2p6

Period 3 3s → 3p blocks 3s23p2 (at Si)

Writing all the sublevel blocks in order gives the

following:

1s22s22p63s23p2

70


Writing Electron Configurations

Using the periodic table, write the electron configuration

for manganese.

Solution:

Period 1 1s block 1s2

Period 2 2s → 2p block 2s2 2p6

Period 3 3s → 3p block 3s2 3p6

Period 4 4s → 3d block 4s2 3d5 (at Mn)

Writing all the sublevel blocks in order gives the

following:

1s22s22p63s23p64s23d5

71




Karen Timberlake

3.8

Trends in Periodic

Table Properties




Valence Electrons

The valence electrons

determine the chemical properties of the elements.

are the electrons in the outermost, highest energy

level.

are related to the group number of the element.

Example: Phosphorus has 5 valence electrons.

5 valence

electrons

P Group 5A(15) 1s22s22p63s23p3

73


Groups and Valence Electrons

All the elements in a group have the same number of

valence electrons.

Example: Elements in Group 2A (2) have two (2)

valence electrons.

Be 1s22s2

Mg 1s22s22p63s2

Ca [Ar]4s2

Sr [Kr]5s2

74


Periodic Table and

Valence Electrons

75


Electron-Dot Symbols

76

An electron-dot symbol

indicates valence electrons

as dots around the symbol of

the element.

of Mg shows two valence

electrons as single dots on the

sides of the symbol Mg.

Mg

Mg Mg Mg Mg


Writing Electron-Dot Symbols

77

The electron-dot symbols for

Groups 1A (1) to 4A (14) use single dots:

Groups 5A (15) to 7A (17) use pairs and single dots:

Na Mg Al C

P O Cl


Groups and Electron-Dot Symbols

In a group, all the electron-dot symbols have the same

number of valence electrons (dots).

Example: Atoms of elements in Group 2A (2) each have

2 valence electrons.

Group 2A (2)

· Be ·

· Mg ·

· Ca ·

· Sr ·

· Ba ·

78


Atomic Size

79

Atomic size

is described using the atomic radius.

is the distance from the nucleus to the valence

electrons.

increases going down a group.

decreases going across a period from left to right.


Atomic Radius

80


Ionization Energy

Ionization energy

is the energy it takes to remove a valence electron

from an atom in the gaseous state.

Na(g) + Energy (ionization) Na+(g) + e–

decreases down a group, increasing across the periodic table from left to right.

81


Ionization Energy and Valence

Electrons

82


Ionization Energy

83

The ionization

energies of

metals are low.

nonmetals are high.


Metallic Character

The metallic character increases when an element can

lose its valence electrons more easily, it

increases down a group where electrons are easier

to remove.

decreases across the period because electrons are

harder to remove.

84


Metallic Character

85


Periodic Table Trend Summary

86

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