Friday 12/04/15

Objectives Understand the general trends in atomic

properties in the periodic table Understand the nature of bonds and their

relationship to electronegativity

Reading the periodic table

Groups or families – vertical columns

Periods – horizontal rows

Term Effective Nuclear Charge (ENC) 1) The net charge that pulls on the valence electrons in an

atom. The greater the effective core charge, the greater the pull. It is determined by subtracting the number of core electrons from the number of protons in the nucleus For example: Magnesium

Periodic Trends

Atomic radius The distance from the center of an atoms nucleus to it’s outermost electron

Measure of atomic size

Periodic Trends Graph the first 20 elements. What is the trend down a group? Across a Period?

Atomic radius

Periodic Trends

Atomic Radius Group Trend

Increases down a group More energy levels or quantum levels (or “shell”) as

you go down a group – atomic radius increases Period Trend

Decreases across a period All electrons in the same energy level. Increased # of

protons holds them closer to nucleus. Increase in Effective Core (Nuclear) Charge (ECC)

Calculate ECC for elements in period 2

Table of Table of Atomic Atomic RadiiRadii

Period Period Trend:Trend:Atomic Atomic RadiusRadius

Periodic Trends Worksheet

Work with a table partner to answer questions:1,5a, 6, 8,

Periodic Trends

Ionic Size

Size of an atom when electrons are added or removed.

Electrons removed atom becomes smaller.

Electrons added atoms become larger

Why? Electron-Electron

Repulsion

Ionic SizeIonic Size

CationsCations

Positively charged ions formed when an atom of a metal loses one or more electrons Smaller than the corresponding atom

AnionsAnions

Negatively charged ions formed when nonmetallic atoms gain one or more electrons

Larger than the corresponding atom

Periodic TrendsGraph the first 20 elements. What is the trend down a group? Across a Period? Ionic Size (label P.T.)

Table of Table of Ion Ion SizesSizes

Ionic Size

Group Trend Increases down a group More energy levels as you go down a group – ionic

size increases Period Trend

Decreases as atoms lose more electrons Increases dramatically as atoms start gaining

electrons, decreases as atoms gain fewer electrons.

Periodic Trends

Ionization Energy Energy needed to remove one of the electrons on an atom’s outer shell.

How strongly does an atom hold it’s outermost electron.

Periodic Trends Graph the first 20 elements. What is the trend down a group? Across a Period?

Ionization Energy

Ionization Energy Group Trends

Increases up a group.The closer outer shell electrons are to the

nucleus the harder they are to remove. Period Trend

Increases across a period.The more electrons in the outer shell the harder

it is to remove one. Increase in Effective Core Charge (ECC)

Periodic Trend:Periodic Trend:Ionization EnergyIonization Energy

Periodic Trends Worksheet

Work with a table partner to answer questions:3, 5b, 7, 9

Periodic Trends

Electronegativity Is a measure of the level of attraction (pull) an atom exerts on the electrons of another atom.

Ability of an atom to attract electrons

Which elements want to gain electrons the most?

Periodic TrendsGraph the first 20 elements. What is the trend down a group? Across a Period? Electronegativity

Periodic Table of ElectronegativitiesPeriodic Table of Electronegativities

Electronegativity Group Trend

Increases up a groupAs radius decreases, electrons are closer to the

nucleus (decrease in number of electron shells) Period Trend

Increases across a periodThe more electrons in the outer shell (up to 7)

the more the atom wants to attract electrons Exception: Trend does not apply to Noble Gases

Increase in Effective Core Charge (ECC)

Periodic Trend:Periodic Trend:ElectronegativityElectronegativity

Periodic Trends Worksheet

Work with a table partner to answer questions:2, 4, 5c, 10, 11

Summary of Summary of Periodic TrendsPeriodic Trends

Practice

1. Se and Br1. Smallest atom

2. Lowest Ionization Energy

2. P, S, Se1. Largest atom

2. Highest Ionization Energy

3. Cl, Cl1-, Br, Br1-

1. Largest ionic size

4. Mg, Mg2+, Na, Na1+

1. Smallest ionic size

Atomic Properties DefinitionsFor Quiz – Monday Effective Nuclear Charge:

It is the net charge that pulls on the valence electrons in an atom.

The greater the effective nuclear charge, the greater the pull.

It is determined by subtracting the number of core electrons from the number of protons in the nucleus

Valence Electrons Are found in the outermost, valence, electron shell

(Bohr model) of the atom Core electrons

occupy all of the inner electron, core, shells

Atomic Properties Definitions Ionization Energy:

Energy needed to remove an electron from an atom or molecule. The higher the effective core charge and lower the number of

electrons shells, the greater the ionization energy Atomic size

How big (e.g., radius) an atom is Atomic radius is measured from the center of the nucleus to the

valence electron shell. The higher the effective core charge and lower the number of

electron shells, the smaller the atom. Electronegativity

Measure of the level of attraction (pull) an atom exerts on the electrons of another atom.

The higher the effective core charge and lower the number of electron shells, the greater the electronegativity

Homework

Read pages:327-331

Answer questions: Pg 336 69-78

343-345 Answer questions: Pgs 374-375, 7-20

Due 12/09

Periodic TableObjective: Students know how to use the periodic table to identify alkali metals, alkaline earth metals, transition metals, metals, semimetals (metalloids), nonmetals, halogens and noble gases.

Alkali Metals All alkali metals have 1

valence electron They are very reactive Reactivity of these elements

increases down the groupAlkali metals:

http://video.google.com/videoplay?docid=-2134266654801392897#

Potassium, K reacts with water and must be stored in kerosene

Alkaline Earth Metals All alkaline earth metals have 2 valence

electrons Alkaline earth metals are less reactive than

alkali metals The word “alkaline” means “basic”

common bases include salts of the metals Ca(OH)2

Mg(OH)2

Properties of MetalsProperties of Metals Metals are good conductors of heat and electricity

Metals are malleable

Metals are ductile

Metals have high tensile strength

Metals have luster

Transition Transition MetalsMetals

Copper, Cu, is a relatively soft metal, and a very good electrical conductor.

Mercury, Hg, is the only metal that exists as a liquid at room temperature

Properties of Metalloids They have properties of both metals and nonmetals.Metalloids are more brittle than metals, less brittle than most nonmetallic solids Metalloids are semiconductors of electricity Some metalloids possess metallic luster

Silicon, Si – A MetalloidSilicon, Si – A Metalloid Silicon has metallic luster Silicon is brittle like a nonmetal Silicon is a semiconductor of electricity

Other metalloids include:

Boron, B Germanium, Ge Arsenic, As Antimony, Sb Tellurium, Te

NonmetalsNonmetals Nonmetals are poor conductors of heat and electricity Nonmetals tend to be brittle Many nonmetals are gases at room temperature

Carbon, the graphite in “pencil lead” is a great example of a nonmetallic element.

Examples of NonmetalsExamples of Nonmetals

Sulfur, S, was once known as “brimstone”

Microspheres of phosphorus, P, a reactive nonmetal

Graphite is not the only pure form of carbon, C. Diamond is also carbon; the color comes from impurities caught within the crystal structure

HalogensHalogens Halogens all have 7 valence electrons

Halogens in their pure form are diatomic molecules (F2, Cl2, Br2, and I2)

Chlorine is a yellow-green poisonous gas

Noble GasesNoble Gases

Noble gases have 8 valence electrons (except helium, which has only 2)

•they are chemically unreactive

• Colorless, odorless and unreactive; they were among the last of the natural elements to be discovered