Unit 3: Periodic Tablehttp://ed.ted.com/periodic-videos

How did chemists begin to organize known elements?

Cu, Au and Ag have been known for thousands of years…. but by 1700 only 13 elements had been identified.

In 1829, Dobereiner first used chemical and physical properties to sort elements into groups (triads).

Mendeleev’s Periodic Table

In 1869, Mendeleev proposed arranging elements by atomic mass and repeating chemical and physical properties.

Mendeleev left gaps on his PT and predicted new elements would be discovered to fill the spaces.

Moseley’s Periodic TableIn 1913, Moseley re-arranged elements by atomic number. Moseley used x-rays to determine the exact atomic number of each element establishing Periodic Law. (physical and chemical properties of elements repeat when elements arranged by increasing atomic number-elements with similar properties appear periodically)

The gaps left in Moseley’s tables had atomic numbers of 43, 61, 72, 75, 86, 87, and 91. These elements are all radioactive!

Organization of the Periodic Table

Group – vertical columns; elements in the same group have same # of electrons in their valence shell and share common characteristics.

Period – horizontal rows; the filling of each energy level with electrons corresponds to a row on the table; do not share common characteristics.

The 3 broad classes of elements

Metals are to the left of the zig-zag line (except H)

Non-metals are to the right of the zig-zag line

Elements touching the line are called “metalloids” (except Al)

Properties of Metals

Shiny

Malleable (can be bent or hammered flat)

Ductile (can be drawn into wire)

Good conductors of heat and electricity

Solids at room temperature (except for ________)

Properties of Nonmetals

Dull

Brittle (nonmalleable)

Poor conductors of heat and electricity

Gases, liquids, or low-melting-point solids

Properties of Metalloids

B, Si, Ge, As, Sb, Te, Po

Have properties of metals and nonmetals

Commonly used in electronics as a semiconductor (ex: Si wafers)

Alkali Metals

Group 1

One valence electron

Very reactive – why?

What is the one exception in group 1 that is not an alkali metal?

Alkaline Earth Metals

Group 2

Two valence electrons

Less reactive than group 1

How to remember names?

Alkali – one word, group 1 Alkaline Earth – two words, group 2

Transition Metals

Groups 3-12

Do not follow patterns as well as groups 1, 2, and 13-18

# of valence electrons harder to predict

Halogens

Group 17

Seven valence electrons

Very reactive –why?

Noble Gases

Group 18

Octet of valence electrons. (Full valence shell)

Inert – unreactive

Don’t form ions

Lanthanide and Actinide Series

Also called Rare Earth Elements

The Lan. Series is part of Period 6The Act. Series is part of Period 7

PERIODIC TRENDS

Periodic Trends

Because of the established Periodic Law, scientists began to notice tendencies of certain elemental characteristics to increase or decrease along a row or column of the periodic table of elements. These tendencies are called Periodic Trends.

4 factors that cause the trends

1. Nuclear Pull (Z) – the number of protons

The protons pull on the outer electrons. The more protons, the more pull exerted by the nucleus on the outer electrons.

2. Electron repulsion –size of e- cloud

The more electrons in

an atom’s electron

cloud, the more they

are pushed away from

each other (due to having

the same charge), making

a bigger cloud.

4 factors that cause the trends

3. Shielding electrons –all inner e- shield the valence electrons from nuclear pull

4 factors that cause the trends

4. Zeff – the “effective” nuclear pull on outer electrons. The nuclear pull taking into account the shielding electrons which are taking most of the force.

Zeff= # protons - # non-valence electrons

Which element has more effective nuclear pull?

Fluorine MagnesiumOR

4 factors that cause the trends

Atomic Radius Trend

• Increases down a column because the valence electrons are in a farther energy level (higher energy)

• Decreases across a period because the nuclear pull is increasing and pulling the energy levels in.

Atomic Size DECREASES

Ato

mic

Siz

e I

NC

REA

ES

Atomic Radius Trend

Using the four factors that determine periodic trends, explain the sizes of the atoms.

IONS

• An atom or group of atoms that has a positive or a negative charge

– Atoms are electrically neutral (protons=electrons)

– Atoms can lose electrons to form cations (+ charge)

– Atoms can gain electrons to form anions (- charge)

Ion SizeMetal ions (cations) are smaller than their atoms

– metal ions LOSE electrons, causing less electron repulsion, and smaller size.

Nonmetal ions (anions) are larger than their atom

– Nonmetal ions GAIN electrons, causing more electron repulsion, and larger size.

Which of the following elements has the largest atomic radii?

K or Rb

If the two elements were to form ions, which element would have the largest ion size?

K or Rb

Ionization Energy

• The energy needed to remove an electron from an atom.

• The greater the ionization energy, the more difficult it is to remove an electron.

Ionization Energy

– Decreases down a group because there are more shielding electrons, so it takes LESS ENERGY to “steal” an electron.

– Increases across a period because the nuclear pull on those electrons is increased with no extra shielding, so it takes MORE ENERGY to get the electrons away.

Ionization Energy INCREASES

Ion

izat

ion

En

erg

y D

ECR

EASE

S

Using the four factors that determine periodic trends, explain why it takes more energy to remove an electron from a fluorine atom vsan oxygen atom.

Electronegativity

• The ability of an atom to take an electron from another atom. Decreases down a group because there are moreelectrons to shield the nucleus (which does thepulling)Increases across a period because of increased Z(nuclear pull)

ElectonegativityINCREASES

Ele

ctro

neg

ativ

ity

DEC

REA

SES

Electron Affinity

• The energy change that occurs when an atom acquires an electron.

• If an atom begins as very unstable (think chaos, lots of energy) and becomes stable by gaining an electron, there was a LARGE ENERGY CHANGE

• Trend correlates with electronegativity, because both involve ability to take electrons

• Decreases down a group and increases across a period

Also important: the most reactive corners of PTare lower left (francium) and upper right (fluorine).Noble gases not included- why?

Using the four factors that determine periodic trends, explain why electrons are more attractive to non-metals than metals.

1. Which is the smallest atom? Explain.Na Li Be

2. Which has the highest electronegativity? Explain.As Sn S

3. In the following pairs, which has the larger atomic radius? Explain.

Mg or Ba Cu or Cu2+ S or S2-

4. In the following pairs, which has the higher ionization energy? Explain.

Li or Cs Ca or Br

Trend What is it?Trend across a

periodTrend across a

groupWhat causes

trend?

Atomic Radius

Ionization Energy

Electronegativity