chapter 4ion.chem.usu.edu/.../lectureslides/ch4students.pdf · general, organic, and biological...

Lecture Presentation

Chapter 4

General, Organic, and Biological Chemistry: Structures of Life, 5/e

Karen C. Timberlake

© 2016 Pearson Education, Inc.

▪ 4.1 – Elements and Symbols

▪ 4.2 – The Periodic Table

▪ 4.3 – The Atom

▪ 4.4 – Atomic Number and Mass Number

▪ 4.5 – Isotopes and Atomic Mass

▪ 4.6 – Electron Energy Levels

▪ 4.7 – Electron Configurations

▪ 4.8 – Trends in Periodic Properties

Goal : Given its name, write its correct symbol; from the symbol. Write the correct name.


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Elements:

▪ are pure substances from which all other things are built.

▪ are listed on the periodic table.


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Element names come from planets, mythological figures,

minerals, colors, geographic locations, and famous people.


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Chemical symbols

• represent the names of the

elements.

• consist of one to two letters

and start with a capital letter.

One-Letter Symbols Two-Letter Symbols

C carbon Co cobalt

N nitrogen Ca calcium

F fluorine Al aluminum O oxygen Mg magnesium

Carbon, C


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Write the correct chemical symbols for each of the following elements:

A. iodine

B. iron

C. magnesium

D. zinc

E. nitrogen


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Give the names of the elements with the following symbols:

A. P

B. Al

C. Mn

D. H

E. K


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▪ Mercury (Hg)

▪ Is a silvery, shiny element that is liquid at room temperature.

▪ Can enter the body by:

▪ mercury vapor inhalation

▪ contact with the skin

▪ ingestion of water or food contaminated with mercury

▪ Once mercury has entered the body, it destroys proteins, disrupts cell function. Long-term exposure can

▪ Damage the brain and kidneys

▪ Cause mental retardation

▪ Decrease physical development


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Mercury contamination comes from

• industrial wastes.

• fish and seafood.

• batteries.

• compact fluorescent bulbs.

Fish absorb mercury.

Big fish eat lots of

small fish, end up with

more mercury.


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Goal : Use the periodic table to identify the group and the period of an element; identify the element as a metal, a nonmetal, or a metalloid.


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The Periodic Table

The periodic table organizes 118 elements into groups with similar properties and places them in order of increasing atomic mass.


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Groups and Periods

In the periodic table,

▪ elements are arranged according to properties.

Vertical columns represent groups of elements

and horizontal rows represent periods.


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Groups and Periods

Group numbers are written at the top of each vertical column (1-18).

Periods are numbered 1-7


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Metals, Nonmetals, and Metalloids

The heavy zigzag line separates metals and nonmetals.

• Metals are located to the left.

• Nonmetals are located to the right.

• Metalloids are located along the heavy zigzag line.

Metals

▪ Solid at room temp▪ Exception: mercury

▪ Shiny!

▪ Ductile (shaped into wires)

▪ Malleable (hammered flat into sheets)

▪ Good conductors

▪ High melting points

Nonmetals

▪ Dull (not shiny )

▪ brittle

▪ Not ductile

▪ Not malleable

▪ Poor conductors

▪ Low melting points

▪ Low densities (many are gasses at room temp.)

Silver

(Metal)

Antimony

(metalloid)

Sulfur

(nonmetal)

Metalloids have a combination of metal and nonmetal properties.


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Characteristics of Metalloids

Metalloids, located along the heavy zigzag line on the periodic table

exhibit properties of metals and nonmetals.

• are better conductors than nonmetals but not as good as metals.

• are used as semiconductors and insulators, because they can be modified to function as conductors or insulators.

Silver

(Metal)

Antimony

(metalloid)

Sulfur

(nonmetal)


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Study Check

Identify each of the following elements as a metal, a nonmetal, or a metalloid:

A. sodium

B. chlorine

C. silicon

D. iron

E. carbon

Study CheckList all of the elements that match the

description.

A. metals in Group 14

Sn, Pb, C, Si, Ge

B. nonmetals in Group 15

Bi, N, P, As, Sb

C. metalloids in Group 14

C, Si, Ge, Sn, Pb


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Group Names


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Alkali Metals

Group 1, the alkali metals, includes the following:

• lithium (Li)

• sodium (Na)

• potassium (K)

• rubidium (Rb)

• cesium (Cs)

• Francium (Fr)

▪ Soft

▪ Shiny

▪ Good conductors

▪ Low melting points

▪ React vigorously with water!


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Alkaline Earth Metals

Group 2 elements, the alkaline earth metals, are shiny but not as reactive as Group 1A metals. They include the following:

• beryllium (Be)

• magnesium (Mg)

• calcium (Ca)

• strontium (Sr)

• barium (Ba)

• radium (Ra)

Strontium gives the red color in

fireworks.


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Halogens

Group 17, the halogens, includes the following:

• fluorine (F)

• chlorine (Cl)

• bromine (Br)

• iodine (I)

• astatine (At)


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Noble Gases

Group 18, the noble gases, include:

▪ Helium (He)

▪ Neon (Ne)

▪ Argon (Ar)

▪ Krypton (Kr)

▪ Xenon (Xe)

▪ Radon (Rn)


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Study Check

Identify the element described by each of the following groups and periods:

1. Group 17, Period 4




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FYI: Elements Essential to Health

Of all the elements, 23 are essential for the well-being and survival of the human body.


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Goal : Describe the electrical charge and location in an atom for a proton, a neutron, and an electron


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An atom is the smallest particle of an element that retains the characteristics of that element.

Aluminum foil contains atoms of aluminum.

The Atom


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Atoms contain the following

subatomic particles:

• Protons

• Neutrons

• Electrons

Atoms are made of subatomic particles


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Atoms contain the following

subatomic particles:

Atoms are made of subatomic particles

• protons have a positive (+) charge

• neutrons have no charge (neutral)

• electrons have a negative (-) charge


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An atom consists of

▪ a nucleus, located in the center of the atom, that contains protonsand neutrons and represents most of the mass of an atom.

▪ electrons that occupy a large, empty space around the nucleus.

Structure of the Atom

Protons – positive charge

Neutrons – neutral

Electrons – negative charge


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Structure of the Atom

In an atom, the protons and neutrons that make up almost all the mass are

packed into the tiny volume of the nucleus. The rapidly moving electrons

(negative charge) surround the nucleus and account for the large volume

of the atom.


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In Dalton’s atomic theory, atoms

• are tiny particles of matter.

• of an element are similar to each other and different from those of other elements.

• of two or more different elements combine to form compounds.

• are rearranged to form new combinations in a chemical reaction.

Atoms are never created or destroyed during a chemical reaction.

Dalton’s Atomic Theory


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Because the mass of subatomic particles are so small,

• chemists use a very small unit of mass called the atomic mass unit (amu).

• 1 amu has a mass equal to 1/12 of the mass of the carbon-12 atom that contains six protons and six neutrons.

• 1 amu = 1 Dalton (Da) in biology.

• 1 amu = 1.66 x 10-27kg

• Electrons have such a small mass that they are not included in the mass of an atom.

Mass of the Atom


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▪ If a proton has a mass of 1.67 x 10-27 kg, what is its mass in amu?


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Protons and neutrons have a very small mass.

Electrons are 1800 times smaller than protons and neutrons.

Subatomic Particles in the Atom


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Which of the following subatomic particles fits each of the

descriptions below?

protons, neutrons, or electrons

A. found outside the nucleus

B. have a positive charge

C. have mass but no charge

Study Check


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Goal : Given the atomic number and the mass number of an atom, state the number of protons, neutrons, and electrons.


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Atomic Number

Each element is assigned an atomic number.

The atomic number is equal to the number of protons in an atom (and is always a whole number).

• The same for all atoms of an element

• Appears above the symbol of an element in the periodic table.


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Atomic Number - examples

Atomic number = number of protons

▪ the atomic number of H is 1; every H atom has one proton.

▪ the atomic number of C is 6; every C atom has six protons.

▪ the atomic number of Cu is 29; every Cu atom has 29 protons.


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Atomic Number - example

All atoms of lithium (left) contain three protons, and all atoms of carbon (right) contain six protons.


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Atoms are Neutral

For neutral atoms, the net charge is zero.

number of protons = number of electrons

Aluminum has 13 protons and 13 electrons. The net (overall)charge is zero.

13 protons (13+) + 13 electrons (13–) = 0


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Study Check

Use the periodic table to fill in the atomic number, number of protons, and number of electrons for each of the following elements:

Element Atomic

Number

Protons Electrons

N

Zn

S


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Mass Number

The mass number

• represents the number of particles in the nucleus.

• is equal to the number of protons + the number of neutrons.

• is always a whole number.

• does not appear in the periodic table.


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Study Tips

Number of protons = atomic number

Number of protons + neutrons = mass number

Number of neutrons = mass number – atomic number

Charge of atom = Number of electrons + Number of protons

Number of electrons = charge – number of protons


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Study Check

An atom of lead (Pb) has a mass number of 207.

A. How many protons are in the nucleus?

B. How many neutrons are in the nucleus?

C. How many electrons are in the atom?


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Study Check

An atom of titanium (Ti) has a mass number of 44.

A. How many protons are in the nucleus?

B. How many neutrons are in the nucleus?

C. How many electrons are in the atom?


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Goal : Determine the number or protons, electrons, and neutrons in one or more of the isotopes of an element; calculate the atomic mass of an element using the percent abundance and mass of its naturally occurring isotopes.


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Isotopes

Isotopes

• are atoms of the same element.

• have the same number of protons but different numbers

of neutrons.

• (different mass numbers)

• can be distinguished by their atomic symbols.


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Atomic Symbols: Subatomic Particles

Given the atomic symbols, determine the number of protons, neutrons, and electrons.

8

16O P31

15 30

65 Zn


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Study Check

Naturally occurring carbon consists of three isotopes: 12C, 13C, and 14C. State the number of protons, neutrons, and electrons in each of the three isotopes.


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Study Check

Write the atomic symbols for atoms with the following subatomic particles:

A. 8 protons 8 neutrons 8 electrons

B. 17 protons 20 neutrons 17 electrons

C. 47 protons 60 neutrons 47 electrons


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Study Check

1. Which of the pairs below are isotopes of the same element?

8

15 X

7

15 X

6

12 X

6

14 X

7

15 X

8

16 X

2. Which of the pairs below have the same number of neutrons?

A B C


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Naturally Occurring Isotopes

Most elements have several isotopes that occur in nature (vs. made in a lab.)

These isotopes are called “naturally occurring.”

Chlorine has 3 isotopes: 35Cl 36Cl 37Cl

Only 35Cl and 37Cl happen naturally.

36Cl has to be made in a lab.

So chlorine has 2 naturally occurring isotopes.


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Atomic Mass

Atomic mass is the

▪ weighted average of all naturally occurring isotopes of that element.

▪ number on the periodic table below the chemical symbol.

Chlorine, with two naturally occurring

isotopes, has an atomic mass of 35.45

amu.


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Isotopes of Magnesium

Magnesium, with three naturally occurring isotopes, has an atomic mass of 24.31 amu.


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Isotopes of Magnesium

Magnesium’s atomic mass: 24.31 amu.


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Atomic Mass of Some Elements


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Calculating Atomic Mass

To calculate atomic mass,

▪ use the experimental percent abundance of each isotope of the element.

▪ multiply the percent abundance (divided by 100) by the atomic mass of that isotope.

▪ sum the total mass of all isotopes.


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Calculating Atomic Mass


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Study Check

Gallium is an element found in lasers used in CD players. In a sample of gallium, there is

60.10% of 69Ga atoms (atomic mass 68.926)

39.90% of 71Ga atoms (atomic mass 70.925)

What is the atomic mass of gallium?


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Goal: Describe the energy levels, sublevels, and orbitals for the electrons in an atom.


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Electromagnetic Radiation

We experience electromagnetic radiation in different

forms, such as light, the colors of a rainbow, or X-rays.


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Electromagnetic Radiation

Light and other electromagnetic radiation consists of energy particles that move as waves of energy.

• The distance between the peaks of waves is called the wavelength.

• High-energy radiation has shorter wavelengths.

• Low-energy radiation has longer wavelengths.

Light is pure energy!All electromagnetic radiation travels

at the speed of light! (3 x 108 m/s!)


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Electromagnetic Spectrum

▪ The Electromagnetic Spectrum above shows all types electromagnetic waves. Many of which we use every day!


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▪ A rainbow appears when sunlight reflects through a raindrop and is separated into its different EM radiation.

The same thing happens when you shine sunlight through a glass prism!


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Atomic Spectrum

They found that each element had it’s own “fingerprint” based on which light it produced. We now call it an element’s atomic spectrum

Scientists found that when they used a different light source than the sun, they didn’t always get a full rainbow!


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▪ The lines of color in an atomic spectrum are caused by the behavior of that element’s electrons.

▪ Electrons occupy specific areas around the atom they are a part of.

Each electron has a specific energy level which corresponds to the

“ring” the electron resides in.

Bohr model and bookcase example


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Electrons and Energy Levels

• Electrons with the same

energy are grouped in the

same energy level.

• Energy levels are assigned

values called principal

quantum numbers (n),

(n = 1, n = 2, …).

An electron can have only the

energy of one of the energy

levels in an atom.


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Electron Energy Levels

In an atom, each electron has a specific energy, known as its energy level, which

▪ is assigned principal quantum numbers (n) = (n = 1, n = 2, …).

▪ increases in energy as the value of n increases and electrons are farther away from the nucleus.

The energy of an electron is quantized—electrons can have only specific energy values.

Bohr model


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Changes in Electron Energy Level

▪ Electrons move to a higher energy level when they absorb energy.

▪ When electrons fall back to a lower energy level, light is emitted.

▪ The energy emitted or absorbed is equal to the differences between the two energy levels.


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Sublevels

It is the arrangement of electrons that determines the physical and chemical properties of an element.

• Each energy level consists of one or more sublevels.

Bookshelf with sublevels

•The number of sublevels in an energy level is

equal to the principal quantum number n of

that energy level.

•The sublevels are identified as s, p, d, and f.

•The order of sublevels in an energy level is

s < p < d < f


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Sublevels

Up to 2 electrons can fit in each “box” (orbital)


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s Orbitals

The location of an electron is described in terms of probability.

• Orbitals are a three-dimensional volume in which electrons have the highest probability of being found.

• The s orbitals are spheres.

(a) The electron cloud of an s orbital

represents the highest probability of

finding an s electron. (b) The s orbitals

are shown as spheres. The sizes of the

s orbitals increase because they

contain electrons at higher energy

levels.


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p Orbitals

There are three p orbitals, starting with n = 2.

• Each p orbital has two lobes, like a balloon tied in the middle, and can hold a maximum of two electrons.

• The three p orbitals are arranged perpendicular to each other along the x, y, and z axes around the nucleus.


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p Orbitals

A p orbital has two regions of high probability, which gives a “dumbbell” shape. (a)

Each p orbital is aligned along a different axis from the other p orbitals. (b) All

three p orbitals are shown around the nucleus.


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d Orbitals

There are five d orbitals, starting with n = 3.

• Four of the five d orbital has four lobes, in the shape of a 4-leaf clover.

• The only difference between them is their location. Shape is identical.

• The 5th d orbital looks like a p orbital with a donut around its middle.


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d Orbitals

Each of the d sublevels contains five d orbitals.

Four of the five d

orbitals consist of

four lobes that are

aligned along or

between different

axes. One d orbital

consists of two lobes

and a doughnut-

shaped ring around

its center.


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Sublevels

Up to 2 electrons can fit in each “box” (orbital)


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▪ Each orbital can hold to two electrons.


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Orbital Capacity and Electron Spin

▪ electrons in the same orbital repel each other.

▪ electrons in the same orbital must have their magnetic spins cancel (they must spin in opposite directions).

We can represent magnetic spins with arrows

Summary: Each orbital can have up to 2 electrons

with one “spin up” and the other “spin down”


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The Pauli exclusion principle states that

▪ each orbital can hold a maximum of two electrons.

▪ electrons in the same orbital repel each other.

▪ electrons in the same orbital must have their magnetic spins cancel (they must spin in opposite directions).


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Number of Electrons in Sublevels

There is a maximum number of electrons that can fill each sublevel.

▪ Each s sublevel has one orbital and can hold a maximum of two electrons.

▪ Each p sublevel has three orbitals and can hold a maximum of six electrons.

▪ Each d sublevel has five orbitals and can hold a maximum of 10 electrons.

▪ Each f sublevel can has 7 orbitals and can hold a maximum of 14 electrons.


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Number of Electrons in Sublevels


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Goal: Draw the orbital diagram and write the electron configuration for an element.


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▪ An orbital diagram shows the placement of electrons in the orbitals in order of increasing energy.

▪ This is much easier than trying to draw all the orbital shapes.

▪ There is 1 s orbital, so thereis 1 box.

▪ There are 3 p orbitals, so there are 3 boxes representing it.

▪ Etc.


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Orbital Diagrams – H and He

We fill an orbital diagram with the number of electrons an atom has.

Element Atomic # # of electrons

Hydrogen (H)

Helium (He)

Begin at the bottom of the chart, putting two electrons in each orbital (box).

Hydrogen Helium

Use an arrow to

represent an

electron.

If there are two

electrons in one

orbital, make one

“up” and one

“down”


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▪ Orbital diagrams can also be written on one line to save space.

▪ They are written in the order of lowest energy to highest.

Hydrogen Helium


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▪ Finally, orbital diagrams can also be written using the orbital names and number of electrons as:

We call this format the electron configuration of that atom or element.


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Electron Configurations

Chemists use a notation called electron configuration to

• indicate the placement of electrons in an atom.

• show how electrons fill energy levels and sublevels in order of increasing energy.

• write an abbreviated form using a noble gas to represent all electrons preceding it.

Electron Configuration for Carbon


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Period 2: Li


Lithium


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Period 2: Be


Beryllium


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Period 2: B, C, and N


Boron

Carbon

Nitrogen

Place one electron in each

p orbital before doubling

up. (same for d and f

orbitals)


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Period 2: O, F, and Ne


Oxygen

Fluorine

Neon

Place one electron in each

p orbital before doubling

up. (same for d and f

orbitals)


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Period 3: Sodium to Argon


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Electron Configurations and the Periodic Table

The electron configurations of elements are related to their positions on the periodic table. Different sections or blocks correspond to sublevels s, p, d, and f.


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Blocks on the Periodic Table

1. The s block contains elements in 1 and 2. This means the final one or two electrons are in the s sublevel.

2. The p block consists of elements in Group 13 to Group 18. There are six p block elements in each period, because three p orbitals can hold a maximum of six electrons.


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Blocks on the Periodic Table

3. The d block, which contains transition elements, first appears after calcium (atomic number 20). There are 10 elements in the d block, because five d orbitals can hold a maximum of 10 electrons.

4. The f block, the inner transition elements, is the two rows of elements at the bottom of the periodic table. There are 14 elements in each f block, because seven f orbitals can hold a maximum of 14 electrons.


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Writing Configurations Using Sublevel Blocks

STEP 1 Locate the element on the periodic table. STEP 2 Write the filled sublevels in order, going across

each period.


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Example: Chlorine


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Electron Configurations: d and f block

Beginning in Period 4, ▪ the 4s sublevel fills before the 3d sublevel, because the 3d

sublevel is slightly lower in energy than the 4s sublevel.▪ the 5s sublevel fills before the 4d sublevel.▪ the 6s sublevel fills before the 5d sublevel.


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Study Check

Use the sublevel blocks on the periodic table to write the electron configuration for selenium.


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d block Sublevel Exceptions

▪ For chromium (Cr), moving one of the 4s electrons to the 3d sublevel adds stability with a half-filled d subshell, and the resulting configuration is 4s13d5.

▪ For copper (Cu), moving one of the 4s electrons to the 3d sublevel adds stability with a filled d subshell, and the resulting configuration is 4s13d10.


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Study Check

Use the sublevel blocks on the periodic table to write the electron configuration for rhenium (Re).


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Study Check

Use the periodic table to give the symbol and name for the element with the electron configuration of

1s22s22p63s23p64s23d7.


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Study Check

Write the electron arrangement for the following elements:

C

Si

O

1s

2s 2p

3s 3p 3d

4s 4p 4d 4f

5s 5p 5d 5f

6s 6p 6d 6f

7s 7p 7d 7f


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Goal: Use the electron configurations of elements to explain the trends in periodic properties.


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Valence Electrons

• Most properties of elements are due to the behavior of

the electrons in the element’s outermost orbitals.

• Valence electrons are the number of electrons in the

outermost energy level.


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Valence Electrons


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Study Check

Using the periodic table, how many valence electrons do each element have?

A. calcium

B. lead


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Lewis Symbols

Lewis symbols represent the valence electrons as dots placed on sides of the symbol for an element.• One to four valence electrons are arranged as single dots.

• Five to eight valence electrons are arranged with at least one pair of electrons around the symbol for the element.

H He B S Kr

Remember, with valence electrons, it doesn’t matter

which row (period) you are in. Only the column (group).


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Study Check

Write the electron-dot (Lewis) symbol for each of the following elements:

Cl NC


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Atomic size

• is determined by the atom’s atomic radius, the distance between the nucleus and the outermost electrons.

Atomic Size


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Atomic Size

The atomic size increases going down a group and right to left across a period.


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▪ Atomic size increases top to bottom down a group

▪ Because each period adds more orbitals which are increasinglyfarther from the nucleus.

▪ Atomic size increases from left to right

▪ Because as you right along a period, you are adding more protons and electrons (but have the same orbitals.) This creates a stronger attraction between the nucleus (protons) and the electrons and it sucks the electrons in tighter. Making the atom smaller.


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Study Check

Given the elements C, N, and Cl,

A. which is the largest atom?


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Ionization Energy

Ionization energy is the energy required to remove one of the outermost (valence) electrons.

• As the distance from the nucleus to the valence electrons increases, the ionization energy decreases.

• The ionization energy is low for metals and high for the nonmetals.


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Ionization Energy

Ionization energy increases up a group and increases going across a period from left to right.


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Study Check

Given the elements C, N, and Cl,

B. which has the highest ionization energy?


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Metallic Character

An element with metallic character is one that loses valence electrons easily (low ionization energy).

Metallic character

• is more prevalent in metals on the left side of the periodic table.

• is less for nonmetals on the right side of the periodic table that do not lose electrons easily.

• decreases going down a group, as electrons are farther away from the nucleus.


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Metallic Character

The metallic character of the s and p block elements increases going down a group and increases going from right to left across a period.


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Summary : Periodic Trends


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Concept Map

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