Chapter 2

The Structure of the Atom and the Periodic Table

Denniston Topping Caret

7th Edition

Copyright The McGraw-Hill Companies, Inc. Permission required for reproduction or display.

2.1 Composition of the Atom

• Atom - the basic structural unit of an element

• The smallest unit of an element that retains the chemical properties of that element

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• Nucleus - small, dense, positively charged region in the center of the atom

- protons - positively charged particles

- neutrons - uncharged particles

Electrons, Protons, and Neutrons

• Atoms consist of three primary particles• electrons• protons• neutrons

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• Electrons are negatively charged particles located outside of the nucleus of an atom

• Protons and electrons have charges that are equal in magnitude but opposite in sign

• A neutral atom that has no electrical charge has the same number of protons and electrons

• Electrons move very rapidly in a relatively large volume of space while the nucleus is small and dense

Mass

number

Atomic number

Charge of particle

Symbol of the atom

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CAZ X

• Atomic number (Z) - the number of protons in the atom

• Mass number (A) - sum of the number of protons and neutrons

Atomic Calculations

number of protons + number of neutrons = mass number

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number of neutrons = mass number - number of protons

number of protons = number of electrons IF positive and negative charges cancel, the atom charge = 0

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Calculate the number of protons, neutrons, and electrons in each of the following:

B115

Fe5526

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Hydrogen(Hydrogen - 1)

Deuterium(Hydrogen - 2)

Tritium(Hydrogen - 3)

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Isotopes of Hydrogen

• Isotopes - atoms of the same element having different masses– contain same number of protons– contain different numbers of neutrons

Isotopes

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• Isotopes of the same element have identical chemical properties

• Some isotopes are radioactive

• Find chlorine on the periodic table

• What is the atomic number of chlorine?

17

• What is the mass given?

35.45

• This is not the mass number of an isotope

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• The atomic mass - the weighted average of the masses of all the isotopes that make up chlorine

• Chlorine consists of chlorine-35 and chlorine-37 in a 3:1 ratio

• Weighted average is an average corrected by the relative amounts of each isotope present in nature

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Calculate the atomic mass of naturally occurring chlorine if 75.77% of chlorine atoms are chlorine-35 and 24.23% of chlorine atoms are chlorine-37

Step 1: convert the percentage to a decimal fraction:

0.7577 chlorine-35

0.2423 chlorine-37

Step 2: multiply the decimal fraction by the mass of that isotope to obtain the isotope contribution to the atomic mass:

For chlorine-35:0.7577 x 35.00 amu = 26.52 amu

For chlorine-370.2423 x 37.00 amu = 8.965 amu

Step 3: sum these partial weights to get the weighted average atomic mass of chlorine:

26.52 amu + 8.965 amu = 35.49 amu

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isotopes– 99.63% nitrogen-14 with a mass of 14.003 amu

– 0.37% nitrogen-15 with a mass of 15.000 amu

• What is the atomic mass of nitrogen?

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Ions and Charges

• Ions - electrically charged particles that result from a gain or loss of one or more electrons by the parent atom

• Cation - positively charged– results from the loss of electrons– 23Na 23Na+ + 1e-

• Anion - negatively charged– results from the gain of electrons– 19F + 1e- 19F-

K3919

-23216S

22412 Mg2.

1 C

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tom Calculating Subatomic Particles

in Ions• How many protons, neutrons, and electrons

are in the following ions?

2.2 Development of Atomic Theory

• Dalton’s Atomic Theory - the first experimentally based theory of atomic structure of the atom

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Postulates of Dalton’s Atomic Theory

1. All matter consists of tiny particles called atoms

2. An atom cannot be created, divided, destroyed, or converted to any other type of atom

3. Atoms of a particular element have identical properties

4. Atoms of different elements have different properties

5. Atoms of different elements combine in simple whole-number ratios to produce compounds (stable aggregates of atoms)

6. Chemical change involves joining, separating, or rearranging atoms

Postulates 1, 4, 5, and 6 are still regarded as true.

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• Electrons were the first subatomic particles to be discovered using the cathode ray tube.

Indicated that the particles were negatively charged.

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Subatomic Particles: Electrons, Protons, and Neutrons

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Evidence for Protons and Neutrons

• Protons were the next particle to be discovered, by Goldstein– Protons have the same size charge but opposite in sign

– A proton is 1,837 times as heavy as an electron

• Neutrons

– Postulated to exist in 1920’s but not demonstrated to exist until 1932

– Almost the same mass as the proton

2.4 The Periodic Law and the Periodic Table

• Dmitri Mendeleev and Lothar Meyer - two scientists working independently developed the precursor to our modern periodic table

• They noticed that as you list elements in order of atomic mass, there is a distinct regular variation of their properties

• Periodic law - the physical and chemical properties of the elements are periodic functions of their atomic numbers

Classification of the Elements2.

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Important Biological Elements2.

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Parts of the Periodic Table

• Period - a horizontal row of elements in the periodic table. They contain 2, 8, 8, 18, 18, and 32 elements

• Group - also called families, and are columns of elements in the periodic table.

• Elements in a particular group or family share many similarities, as in a human family.2.

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• Representative elements - Group A elements

• Transition elements - Group B elements

• Alkali metals - Group IA

• Alkaline earth metals - group IIA

• Halogens - group VIIA

• Noble gases - group VIIIA

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Category Classification of Elements

• Metals - elements that tend to lose electrons during chemical change, forming positive ions

• Nonmetals - a substance whose atoms tend to gain electrons during chemical change, forming negative ions

• Metalloids - have properties intermediate between metals and nonmetals

Classification of Elements Metals

• Metals: – A substance whose atoms tend to lose

electrons during chemical change– Elements found primarily in the left 2/3 of

the periodic table

• Properties:– High thermal and electrical conductivities– High malleability and ductility– Metallic luster– Solid at room temperature

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Classification of Elements Nonmetals

• Nonmetals: – A substance whose atoms may gain

electrons, forming negative ions– Elements found in the right 1/3 of the

periodic table

• Properties:– Brittle– Powdery solids or gases– Opposite of metal properties

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Classification of Elements Metalloids

• Metalloids: – Elements that form a narrow diagonal band

in the periodic table between metals and nonmetals

• Properties are somewhat between those of metals and nonmetals

• Also called semimetals

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Atomic Number and Atomic Mass

• Atomic Number:

– The number of protons in the nucleus of an atom of an element

– Nuclear charge or positive charge from the nucleus

• Most periodic tables give the element symbol, atomic number, and atomic mass

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Element Information in the Periodic Table

20 atomic number

Ca symbol

Calcium name

40.08 atomic mass

Using the Periodic Table

• Identify the group and period to which each of the following belongs:

a. P

b. Cr

c. Element 30

• How many elements are found in period 6?

• How many elements are in group VA?2.

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2.5 Electron Arrangement and the Periodic Table

• The electron arrangement is the primary factor in understanding how atoms join together to form compounds

• Electron configuration - describes the arrangement of electrons in atoms

• Valence electrons - outermost electrons– The electrons involved in chemical bonding

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• The number of valence electrons is the group number for the representative elements

• The period number gives the energy level (n) of the valence shell for all elements

Valence Electrons and Energy Level

• How many valence electrons does Fluorine have?

– 7 valence electrons

• What is the energy level of these electrons?

– Energy level is n = 2

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Electron Arrangement by Energy Level

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Valence Electrons - Detail• What is the total number of electrons in

fluorine?

– Atomic number = 9

– 9 protons and 9 electrons

• 7 electrons in the valence shell, (n = 2 energy level), so where are the other two electrons?– In n = 1 energy level

– Level n = 1 holds only two electrons

Determining Electron ArrangementList the total number of electrons, total number ofvalence electrons, and energy level of the valenceelectrons for silicon.

1. Find silicon in the periodic table• Group IVA • Period 3• Atomic number = 14

2. Atomic number = number of electrons in an atom

• Silicon has 14 electrons

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Determining Electron Arrangement #2List the total number of electrons, total number of valence electrons, and energy level of the valence electrons for silicon.

3. As silicon is in Group IV, only 4 of its 14 electrons are valence electrons

• Group IVA = number of valence electrons

4. Energy levels:• n = 1 holds 2 electrons• n = 2 holds 8 electrons (total of 10)

• n = 3 holds remaining 4 electrons (total = 14)

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Determining Electron ArrangementPractice

List the total number of electrons, total number of valence electrons, and energy level of the valence electrons for:

• Na

• Mg

• S

• Cl

• Ar

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PRINCIPAL ENERGY LEVELS

• n = 1, 2, 3, …

• The larger the value of n, the higher the energy level and the farther away from the nucleus the electrons are

• The number of sublevels in a principal energy level is equal to n

– in n = 1, there is one sublevel

– in n = 2, there are two sublevels

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Principal Energy Levels

• The electron capacity of a principal energy level (or total electrons it can hold) is

2(n)2

– n = 1 can hold 2(1)2 = 2 electrons

– n = 2 can hold 2(2)2 = 8 electrons

• How many electrons can be in the n = 3 level?– 2(3)2 = 18

• Compare the formula with periodic table…..

n = 1, 2(1)2 = 2

n = 2, 2(2)2 = 8

n = 3, 2(3)2 = 18

n = 4, 2(4)2 = 32

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Sublevels

• Sublevel: a set of energy-equal orbitals within a principal energy level

• Subshells increase in energy:

s < p < d < f

• Electrons in 3d subshell have more energy than electrons in the 3p subshell

• Specify both the principal energy level and a subshell when describing the location of an electron

Principle energy level (n)

Possible subshells

1 1s

2 2s, 2p

3 3s, 3p, 3d

4 4s, 4p, 4d, 4f

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Orbitals

• Orbital - a specific region of a sublevel containing a maximum of two electrons

• Orbitals are named by their sublevel and principal energy level

– 1s, 2s, 3s, 2p, etc.

• Each type of orbital has a characteristic shape

– s is spherically symmetrical

– p has a shape much like a dumbbell

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• s is spherically symmetrical

• Each p has a shape much like a dumbbell, differing in the direction extending into space

Subshell Number of

orbitals

s 1

p 3

d 5

f 7

• How many electrons can be in the 4d subshell?

•10

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Quantum Mechanical Model• Each orbital within a

sublevel contains a maximum of 2 electrons

• Energy increases as n, shell number increases, but ALSO increases as you move from s to p to d to f sublevels

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Incr

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4p

4d

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••

•• •• ••

•• •• •• •• ••

••••••••••••••

Electron

Orbital

Sublevel

Shell 4

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Electron Spin• Electron configuration - the

arrangement of electrons in atomic orbitals

• Aufbau principle - or building up principle helps determine the electron configuration– Electrons fill the lowest-energy orbital that

is available first– Remember s<p<d<f in energy– When the orbital contains two electrons,

the electrons are said to be paired

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Rules for Writing Electron Configurations

• Obtain the total number of electrons in the atom from the atomic number

• Electrons in atoms occupy the lowest energy orbitals that are available – 1s first

• Each principal energy level, n contains only n sublevels

• Each sublevel is composed of orbitals• No more than 2 electrons in any orbital• Maximum number of electrons in any principal

energy level is 2(n)2

Electron Distribution• This table lists the number of electrons in each

shell for the first 20 elements• Note that 3rd shell stops filling at 8 electrons even though

it could hold more

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Orbital Energy-level Diagram2.

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Writing Electron Configurations

• H– Hydrogen has

only 1 electron– It is in the

lowest energy level & lowest orbital

– Indicate number of electrons with a superscript

– 1s1

• Li– Lithium has 3

electrons– First two have

configuration of Helium – 1s2

– 3rd is in the orbital of lowest energy in n=2

– 1s2 2s1

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• Give the complete electron configuration of each element

– Be

– N

– Na

– Cl

– Ag

The Shell Model and Chemical Properties

• As we explore the model placing electrons in shells, we will see that the pattern which emerges from this placement correlates well with a pattern for various chemical properties

• We will see that all elements in a group have the same number of electrons in their outermost (or valence) shell2.

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Groups Have Similar Chemical Properties and Appearances

• Examples of different elements that have similar properties and are all in group VA– Nitrogen

– Phosphorus

– Arsenic

– Antimony

– Bismuth

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What noble gas configuration is this?

•Neon•Configuration is written: [Ne]3s23p1

Shorthand Electron Configurations

• Uses noble gas symbols to represent the inner shell and the outer shell or valance shell is written after

• Aluminum- full electron configuration is: 1s22s22p63s23p1

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• Remember:

– How many subshells are in each principle energy level?

– There are n subshells in the n principle energy level.

– How many orbitals are in each subshell?

– s has 1, p has 3, d has 5, and f has 7

– How many electrons fit in each orbital?

– 22.5

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Shorthand Electron Configuration Examples

• N

• S

• Ti

• Sn

Use this breakdown of the Periodic Table and you can write the configuration of any element.

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Classification of Elements According to the Type of

Subshells Being Filled

Classification of Elements – by Group

• Representative element: An element in which the distinguishing electron is found in an s or p subshell

• Distinguishing electron: The last or highest-energy electron found in an element

• Transition element: An element in which the distinguishing electron is found in a d subshell

• Inner-transition element: An element in which the distinguishing electron is found in a f subshell

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2.6 The Octet Rule

• The noble gases are extremely stable– Called inert as they don’t readily bond to other

elements

• The stability is due to a full complement of valence electrons in the outermost s and p sublevels:– 2 electrons in the 1s of Helium – the s and p subshells are full in the outermost

shell of the other noble gases (eight electrons)

Octet of Electrons

• Elements in families other than the noble gases are more reactive– Strive to achieve a more stable electron

configuration– Change the number of electrons in the atom to

result in full s and p sublevels

• Stable electron configuration is called the “noble gas” configuration2.6

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The Octet Rule

• Octet rule - elements usually react in such a way as to attain the electron configuration of the noble gas closest to them in the periodic table– Elements on the right side of the table move right to the

next noble gas– Elements on the left side move “backwards” to the

noble gas of the previous row

• Atoms will gain, lose or share electrons in chemical reactions to attain this more stable energy state

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NaSodium atom

11e-, 1 valence e-

[Ne]3s1

Na+ + e-

Sodium ion10e-

[Ne]

Ion Formation and the Octet Rule

• Metallic elements tend to form positively charged ions called cations

• Metals tend to lose all their valence electrons to obtain a configuration of the noble gas

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AlAluminum atom13e-, 3 valence e-

[Ne]3s23p1

Al3+ + 3e-

Aluminum ion10e-

[Ne]

• All atoms of a group lose the same number of electrons

• Resulting ion has the same number of electrons as the nearest (previous) noble gas atom

Ion Formation and the Octet Rule

O + 2e-

Oxygen atom8e-, 6 valence e-

[He]2s22p4

O2-

Oxide ion10e-

[He]2s22p6 or [Ne]

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Isoelectronic• Isoelectronic - atoms of different elements having

the same electron configuration (same number of electrons)

• Nonmetallic elements, located on the right side of the periodic table, tend to form negatively charged ions called anions

• Nonmetals tend to gain electrons so they become isoelectronic with its nearest noble gas neighbor located in the same period to the right

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Using the Octet Rule

• The octet rule is very helpful in predicting the charges of ions in the representative elements

• Transition metals still tend to lose electrons to become cations but predicting the charge is not as easy

• Transition metals often form more than one stable ion– Iron forming Fe2+ and Fe3+ is a common example

Examples Using the Octet Rule

• Give the charge of the most probable ion resulting from these elements– Ca

– Sr

– S

– P

• Which of the following pairs of atoms and ions are isoelectronic?– Cl-, Ar

– Na+, Ne

– Mg2+, Na+

– O2-, F-2.6

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2.7 Trends in the Periodic Table

• Many atomic properties correlate with electronic structure and so also with their position in the periodic table– atomic size– ion size– ionization energy– electron affinity

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leAtomic Size

• The size of an element increases, moving down from top to bottom of a group

• The valence shell is higher in energy and farther from the nucleus traveling down the group

• The size of an element decreases from left to right across a period

• The increase in magnitude of positive charge in nucleus pulls the electrons closer to the nucleus

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leVariation in Size of Atoms

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leCation Size

Cations are smaller than their parent atom• More protons than electrons creates an increased

nuclear charge• Extra protons pull the remaining electrons closer

to the nucleus• Ions with multiple positive charges are even

smaller than the corresponding monopositive ions– Which would be smaller, Fe2+ or Fe3+? Fe3+

• When a cation is formed isoelectronic with a noble gas the valence shell is lost, decreasing the diameter of the ion relative to the parent atom

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leAnion Size

Anions are larger than their parent atom.

• Anions have more electrons than protons

• Excess negative charge reduces the pull of the nucleus on each individual electron

• Ions with multiple negative charges are even larger than the corresponding monopositive ions

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Their Parent Atoms

ionization energy + Na Na+ + e-2.7

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leIonization Energy

• Ionization energy - The energy required to remove an electron from an isolated atom

• The magnitude of ionization energy correlates with the strength of the attractive force between the nucleus and the outermost electron

• The lower the ionization energy, the easier it is to form a cation

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leIonization Energy of Select Elements

• Ionization decreases down a family as the outermost electrons are farther from the nucleus

• Ionization increases across a period because the outermost electrons are more tightly held

• Why would the noble gases be so unreactive?

Br + e– Br– + energy

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• Electron affinity - The energy released when a single electron is added to an isolated atom

• Electron affinity gives information about the ease of anion formation

– Large electron affinity indicates an atom becomes more stable as it forms an anion

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lePeriodic Trends in Electron

Affinity

• Electron affinity generally decreases down a group

• Electron affinity generally increases across a period