karen timberlake chapter 3 - chemistry121 [licensed for...

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General, Organic, and Biological Chemistry

Fourth 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

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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 SymbolsC carbon Co cobaltN nitrogen Ca calciumF fluorine Al aluminum O oxygen Mg magnesium

Symbols of Elements


The Periodic Table

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• 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.

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Periods and Groups

7 © 2013 Pearson Education, Inc. Chapter 3, Section 2

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, elementsinclude 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

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Main Group Elements

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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).


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.

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Fourth Edition

Karen Timberlake

Chapter 3The Atom

© 2013 Pearson Education, Inc.

A Brief History of

Atomic Theory


John Dalton’s Atomic Theory ( circa

1804)

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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

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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.

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Rutherford’s Gold-Foil

Experiment

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

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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.

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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 Levels

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� 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



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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.

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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.


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 massof 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

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Particles in the Atom

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• 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.

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Na

Atomic Number

Symbol


Atomic Number and Protons

Each element has a unique atomic number equal to thenumber 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.


Number of Electrons in an Atom

All atoms of an element are electrically neutral; theyhave

� 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; thenet charge is zero.

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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.



Fourth Edition

Karen Timberlake

3.5Isotopes 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.


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Isotopes and Mass and Atomic Symbols

(Nuclear, or Isotopic, Notation)

Since each isotope of an element has a differentnumber of neutrons, each isotope’s mass number willbe 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 11and a mass number 23 has the following atomic symbol:

mass number atomic number

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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

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Average Atomic Mass

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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

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Most elements have two or more isotopes thatcontribute 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.


Calculating the Atomic Mass for Cl

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To calculate the atomic mass of an element, we need to know the percent abundance of each isotope and its mass.



Fourth Edition

Karen Timberlake

3.6Electron 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 andchemical properties of atoms.

So, we must learn more about the electronic structure of the elements.

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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.


Electron Energy Levels

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Electron energy levels increasein energy and number aselectrons get farther away fromthe 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 subleveln=2 has 2 sublevels – the 2s and the 2p subleveln=3 has 3 sublevels – the 3s, 3p, and 3d subleveln=4 has 4 sublevels – the 4s, 4p, 4d, and 4f sublevel


Sublevels

n= 1 has 1 sublevel -- the 1s subleveln=2 has 2 sublevels – the 2s and the 2p subleveln=3 has 3 sublevels – the 3s, 3p, and 3d subleveln=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.

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Energy Levels , Sublevels and

Orbital types

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� 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.

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Orbitals

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Each electron sublevel consists of orbitals, which

� are regions where there is the highest probabilityof finding an electron.

� have their own unique three-dimensional shape.

� Each can hold up to 2 electrons.


s Orbitals

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We know that s orbitals have a spherical shape, centered around the atom’s nucleus (located at theorigins 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

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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

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There are five d 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.


Electron Capacity in Sublevels

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Fourth Edition

Karen Timberlake

3.7Orbital 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.

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Energy Diagram for Sublevels


Orbital Diagrams

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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 Li1s2

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.


Writing Orbital Diagrams

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The orbital diagram forcarbon 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.

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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


Period 1 Configurations

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



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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.

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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 lessthan the period number.) = d block

� Lanthanides/Actinides (This sublevel is (n-2), 2 lessthan the period number.) = f block


Sublevel Blocks

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Guide to Using Sublevel Blocks


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

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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



Fourth Edition

Karen Timberlake

3.8Trends in Periodic

Table Properties



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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


Groups and Valence Electrons

All the elements in a group have the same number ofvalence electrons.

Example: Elements in Group 2A (2) have two (2) valence electrons.

Be 1s22s2

Mg 1s22s22p63s2

Ca [Ar]4s2

Sr [Kr]5s2

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Periodic Table and

Valence Electrons


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

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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 have2 valence electrons.

Group 2A (2)

· Be ·

· Mg ·

· Ca ·

· Sr ·

· Ba ·

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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

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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.


Ionization Energy and Valence

Electrons

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Ionization Energy

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The ionizationenergies 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.

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Metallic Character


Periodic Table Trend Summary

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