1

3.1.1 Describe the arrangement of elements in the periodic table in order of increasing atomic

number.

Development of the Periodic Table

• Johan Dobereiner Grouped similar elements into groups of 3 (triads) such as chlorine, bromine, and iodine. (1817-1829).

• John Newlands

Found every eighth element (arranged by atomic weight) showed similar properties. Law of Octaves (1863).

• Dmitri Mendeleev Arranged elements by similar properties but left blanks for undiscovered elements (1869).

Dmitri Mendeleev 1834 – 1907

• Russian chemist and teacher

• given the elements he knew about, he organized a “Periodic Table” based on increasing atomic mass (it’s now atomic #)

• he even left empty spaces to be filled in later

At the time the elements gallium and germanium were not known. These are the blank spaces in his periodic table. He

predicted their discovery and estimated their properties.

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IB Topic 3: Periodicity 3.1: The periodic table

• 3.1.1 Describe the arrangement of elements in the periodic table in order of increasing atomic number.

• 3.1.2 Distinguish between the terms group and period.

• 3.1.3 Apply the relationship between the electron arrangement of elements and their position in the periodic table up to Z = 20.

• 3.1.4 Apply the relationship between the number of electrons in the highest occupied energy level for an element and its position in the periodic table.

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Henry Moseley 1887 – 1915

• arranged the elements in increasing atomic numbers (Z)– properties now

recurred periodically

Design of the Table• Groups are the vertical columns.

– elements have similar, but not identical, properties• most important property is that

they have the same # of valence electrons

• valence electrons- electrons in the highest occupied energy level

• all elements have 1,2,3,4,5,6,7, or 8 valence electrons

IB prefers this one.

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3.1.1 Describe the arrangement of elements in the periodic table in order of increasing atomic number.3.1.2 Distinguish between the terms group and period.

Development of the Periodic Table

• Henry Mosley Arranged the elements by increasing atomic number instead of mass (1913)

• Glen Seaborg Discovered the transuranium elements (93-102) and added the actinide and lanthanide series (1945)

Elements arranged by increasing atomic number into • periods (rows) and • groups or families (columns),

which share similar characteristics

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3.1.1 Describe the arrangement of elements in the periodic table in order of increasing atomic

number.

Metals

• Left side of the periodic table (except hydrogen)

• Good conductors of heat and electricity

• Malleable: capable of being hammered into thin sheets

• Ductile: capable of being drawn into wires

• Have luster: are shiny• Typically lose electrons in

chemical reactions

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3.1.1 Describe the arrangement of elements in the periodic table in order of increasing atomic

number.

Metals

• Alkali metals: Group 1 (1A)• Alkaline earth metals: Group 2

(2A)• Transition metals: Group B,

lanthanide & actinide series

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3.1.1 Describe the arrangement of elements in the periodic table in order of increasing atomic

number.

Nonmetals

• Right side of the periodic table • Poor conductors of heat and

electricity• Non-lusterous• Typically gain electrons in

chemical reactions

• Halogens: Group 17 (7A)• Noble gases: Group 18 (0)

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3.1.1 Describe the arrangement of elements in the periodic table in order of increasing atomic

number.

Metalloids

• Between metals and non-metals, along the stair step (except aluminum)

• Have properties of metals and non-metals

• Some are semi-conductors

• Boron (B), Silicon (Si), Germanium (Ge), Arsenic (As), Antimony (Sb), Tellurium (Te), Astatine (At)

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16

ns1

ns2

ns2

np1

ns2

np2

ns2

np3

ns2

np4

ns2

np5

ns2

np6

d1

d5 d10

4f

5f

Ground State Electron Configurations of the Elements

Electron ArrangementElectron Arrangement

Core Electrons: electrons that are in the inner energy levels

Valence Electrons: electrons that are in the outermost (highest) energy level

Group = Sum of electrons in the highest occupied energy level (s + p) = Number of valence electrons

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3.1.4 Apply the relationship between the number of electrons in the highest occupied energy level

for an element and its position in the periodic table.

Arrangement of the Periodic Table

• Valence Electrons: electrons in the outermost (highest) energy level– Group 1 elements have 1 v.e.s – Group 2 elements have 2 v.e.s– Group 3 elements have 3 v.e.s – So on and so forth– Group 8 have 8 v.e. (except

for helium, which has 2)

Lewis Dot-Diagrams/Structures

• valence electrons are represented as dots around the chemical symbol for the element

Na

Cl

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3.1.4 Apply the relationship between the number of electrons in the highest occupied energy level

for an element and its position in the periodic table.

Electron dot diagrams

Group 1A: 1 dot X Group 5A: 5 dots X

Group 2A: 2 dots X Group 6A: 6 dots X

Group 3A: 3 dots X Group 7A: 7 dots X

Group 4A: 4 dots X Group 0: 8 dots (except He) X

Look, they are following my

rule!

Electron Dot Diagram Using the symbol for the element, place dots around the symbol

correspondingto the outer energy level s & p electrons (valence electrons). Will

have fromone to eight dots in the dot diagram.

Draw electron dot diagrams for the following atoms

H Be O Al Ca Zr

H Be O

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Electron Dot Diagram Using the symbol for the element, place dots around the symbol

correspondingto the outer energy level s & p electrons. Will have from one to eight

dots inthe dot diagram.

Draw electron dot diagrams for the following atoms

Al Ca Zr

Al Ca Zr

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2.3.4 Deduce the electron arrangement 2.3.4 Deduce the electron arrangement for atoms and ions.for atoms and ions.

Write electron configuration, orbital filling diagrams, and electron dot

diagrams.

Kr

Tb 26

• B is 1s2 2s2 2p1;– 2 is the outermost energy level – it contains 3 valence electrons, 2

in the 2s and 1 in the 2p• Br is [Ar] 4s2 3d10 4p5

How many valence electrons are present?

• Periods are the horizontal rows– do NOT have similar properties– however, there is a pattern to their properties

as you move across the table that is visible when they react with other elements

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3.1.3 Apply the relationship between the electron arrangement of elements and their position in the

periodic table.

Arrangement of the Periodic Table

• Period = The highest occupied energy level = number of energy levels

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3.1.3 Apply the relationship between the electron arrangement of elements and their position in the

periodic table.

Arrangement of the Periodic Table

• Na = 1s22s22p63s1

• Sodium is in the 3rd period because it has 3 energy levels The highest occupied energy level is 3

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IB Topic 3: Periodicity 3.2: Physical properties

• 3.2.1 Define the terms first ionization energy and electronegativity.

• 3.2.2 Describe and explain the trends in atomic radii, ionic radii, first ionization energies, electronegativities and melting points for the alkali metals (Li Cs) and the halogens (F I).

• 3.2.3 Describe and explain the trends in atomic radii, ionic radii, first ionization energies and electronegativities for elements across period 3.

• 3.2.4 Compare the relative electronegativity values of two or more elements based on their positions in the periodic table.

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Periodic Trend Definitions

• Atomic Radius: half the internuclear distance between two atoms of the same element (pm)

• Ionic radius: the radius of an ion in the crystalline form of a compound (pm)

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Periodic Trend Definitions

• First ionization energy: The energy required to remove one electron from each atom in one mole of gaseous atoms under standard thermodynamic conditions (kJ mol-1)

• Electron Affinity: The energy released when one electron is added to each atom in one mole of gaseous atoms under standard thermodynamic conditions (kJ mol-1)

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Periodic Trend Definitions

• Electronegativity: a measure of the tendency of an atom in a molecule to attract a pair of shared electrons towards itself

• Melting Point: the temperature at which a solid becomes a liquid at a fixed pressure (degrees Kelvin)

Trends in the table IB loves the alkali metals and

the halogens

• many trends are easier to understand if you comprehend the following

• the ability of an atom to “hang on to” or attract its valence electrons is the result of two opposing forces – the attraction between the electron and

the nucleus– the repulsions between the electron in

question and all the other electrons in the atom (often referred to the shielding effect)

– the net resulting force of these two is referred to effective nuclear charge

This is a simple, yet very good picture. Do you understand it?

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3.2.2 Describe and explain the trends in atomic radii, ionic radii, first ionization

energies, electronegativities and melting points

Group 1A: Alkali Metals• Have 1 valence electron• Shiny, silvery, soft metals• React with water & halogens• Oxidize easily (lose electrons)• Reactivity increases down the

group

Group 7A: Halogens • Have 7 valence electrons• Colored gas (F2, Cl2); liquid (Br2);

Solid (I2)

• Oxidizer (gain electrons)• Reactivity decreases down the

group

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3.2.2 Describe and explain the trends in atomic radii, ionic radii, first ionization

energies, electronegativities and melting points

Atomic Radii• The radius of an atom, measured in pm

(picometers)

• Periodic trend (Period 3 Trend)– Atomic size decreases as you move across a period.– The increase in nuclear charge increases the attraction to

the outer shell so the outer energy level progressively becomes closer to the nucleus

• Group trend for Alkali metals & Halogens– Atomic size increases as you move down a group of the

periodic table.– Adding higher energy levels

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Atomic Radii

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3.2.2 Describe and explain the trends in atomic radii, ionic radii, first ionization

energies, electronegativities and melting points

Ionic RadiiThe radius of the ion form of atoms

(cations and anions)

• Positive ions are smaller than their atoms.– Fewer electrons so nucleus attracts remaining electrons more strongly– One fewer energy level since valence electrons removed.

• Negative ions are larger than their atoms– More electrons so nucleus has less attraction for them– Greater electron-electron repulsion

• Periodic trend (Period 3 Trend)– Decrease as you move across a period, then spike and decrease again – This increase in nuclear charge increases the attraction to the outer shell so the

outer energy level progressively becomes closer to the nucleus

• Group trend for Alkali metals & Halogens– Ions get larger down a group– More energy levels are added

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45

3.2.2 Describe and explain the trends in atomic radii, ionic radii, first ionization

energies, electronegativities and melting points

First Ionization EnergiesThe energy required to remove the first electron from a

gaseous atom.Second ionization removes the second electron and so on.

Can be used to predict ionic charges.

• Periodic Trend (Period 3 Trend)– Increases as you move from left to right across a period.– Effect of increasing nuclear charge makes it harder to remove an

electron.

• Group trend for Alkali metals & Halogens– Generally decreases as you move down a group in the periodic table – Since size increases down a group, the outermost electron is farther

away from the nucleus and is easier to remove.

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Filled n=1 shell

Filled n=2 shell

Filled n=3 shellFilled n=4 shell

Filled n=5 shell

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3.2.2 Describe and explain the trends in atomic radii, ionic radii, first ionization

energies, electronegativities and melting points

ElectronegativityTendency for the atoms of the element to attract

electrons when theyare chemically combined with atoms of another

element.Helps predict the type of bonding (ionic/covalent).

• Periodic Trend (Period 3 Trend)– Increases as you move from left to right across a period.– Nonmetals have a greater attraction for electrons than metals &

there is a greater nuclear charge that can attract electrons

• Group trend for Alkali metals & Halogens– Generally decreases as you move down a group in the periodic

table.– For metals, the lower the number the more reactive.– For nonmetals, the higher the number the more reactive.

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ElectronegativityElectronegativity

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3.2.2 Describe and explain the trends in atomic radii, ionic radii, first ionization

energies, electronegativities and melting points

ReactivityThe relative capacity of an atom, molecule or radical to

undergo a chemical reaction with another atom, molecule or radical.

• Don’t worry about the periodic trend!!!

• Group trend for Alkali metals– Increases as you move down group 1 in the periodic table – Since alkali metals are more likely to lose an electron, the ones

with the lowest 1st ionization energy are the most reactive since they require the least amount of energy to lose a valence electron.

• Group trend for Halogens– Decreases as you move down group 7 in the periodic table – Since halogens are more likely to gain an electron, the ones with

the greatest electronegativity are the most reactive since they are most effective at gaining a valence electron.

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3.2.2 Describe and explain the trends in atomic radii, ionic radii, first ionization

energies, electronegativities and melting points

Melting PointsThe temperature at which a crystalline melts depends on the strength of the attractive forces and on the way the particles are packed in the

solid state

• Don’t worry about the periodic trend!!!

• Alkali Metals: Melting point decreases down the group– Li (181 oC) to Cs (29 oC)– As the atoms get larger the forces of attraction between them

decrease due to the type of bonding (metallic)

• Halogens: Melting point increases down the group– F2 (-220 0C) to I2 (114 oC)– Weak attractive forces increase as the molecules get larger

due to the type of bonding (non-polar covalent)

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IB Topic 3: Periodicity 3.3: Chemical properties

• Discuss the similarities and differences in the chemical properties of elements in the same group.

• Discuss the changes in nature, from ionic to covalent and from basic to acidic, of the oxides across period 3.

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3.3.1 Discuss the similarities and differences in the chemical

properties of elements in the same group.

Alkali Metals

React with water &react with many

substances because…

They have the same number of

valence electrons

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3.3.1 Discuss the similarities and differences in the chemical

properties of elements in the same group.

Alkali Metals

2Na(s) + 2H2O(l) 2NaOH (aq) + H2(g)

In the reaction of alkali metals and water, all will:

• move around the surface of the water,

• give off hydrogen gas, • create a basic

solution.

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3.3.1 Discuss the similarities and differences in the chemical

properties of elements in the same group.

Alkali Metals

In the reaction of alkali metals and water, the reactivity will

increase down the group because they get better at getting rid of their valence

electron (the 1st ionization energy

decreases)

So, alkali metals lower down will:

• React more vigorously• React faster

• Give off a flame

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3.3.1 Discuss the similarities and differences in the chemical

properties of elements in the same group.

Alkali Metals

Reaction with halogens

2M(s) + X2 (g) 2MX(s)

where M represents Li,Na,K,Rb, or Cs

Where X represents F,Cl,Br, or I

2Na(s) + Cl2(g) 2NaCl(s)

Reactivity decreases down the group

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3.3.1 Discuss the similarities and differences in the chemical

properties of elements in the same group.Halogens

Halogens are diatomic as gases (two atoms bond together) and called halides when they form ions… These are BrINClHOF

Halogens want to get one electron to fill its outer shell.

Reactivity decreases down the group because electronegativity decreases

Cl2 reacts with Br- and I- Cl2(aq) + 2Br-(aq) 2Cl-(aq) + Br2(l)

Cl2(aq) + 2I-(aq) 2Cl-(aq) + I2(s)

Br2 reacts with I-Br2(aq) + 2I-(aq) 2Br-(aq) + I2(s)

I2 non-reactive with halide ions

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Reactivity of Elements… in action

Alkali Metals: http://www.youtube.com/watch?v=m55kgyApYrY

Halogens:

http://www.youtube.com/watch?v=tk5xwS5bZMA&feature=related

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3.3.2 Discuss the changes in nature, from ionic to covalent and from

basic to acidic, of the oxides across period 3.

Metallic Oxides in Period 3Sodium oxide: Na2O ionic

Magnesium oxide: MgO ionicAluminum oxide: Al2O3 ionic

Metalloid oxide in Period 3Silicon dioxide: SiO2 covalent

Nonmetallic oxides in Period 3Tetraphosphorus decoxide: P4O10 covalent

Sulfur trioxide: SO3 covalent

Dichlorine heptoxide: Cl2O7 covalent

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3.3.2 Discuss the changes in nature, from ionic to covalent and from

basic to acidic, of the oxides across period 3.

Acidic/Basic

Metallic oxides in Period 3 are basicSodium oxide: Na2O + H2O 2 NaOH basicMagnesium oxide: MgO + H2O Mg(OH)2 basicAluminum oxide: Al2O3 + H2O 2 Al(OH)3 amphoteric

Metalloid oxide in Period 3 is acidicSilicon dioxide: SiO2 + H2O H2SiO3 acidic

Nonmetallic oxides in Period 3 are acidicTetraphosphorus decoxide: P4O10 + 6H2O 4H3PO4 acidicSulfur trioxide: SO3 + H2O H2SO4 acidicDichlorine heptoxide: Cl2O7 + H2O 2HClO4 acidicArgon does not form an oxide

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Terms to Know

• Group

• Period

• Alkali metals

• Halogens

• Ionic radius

• Electronegativity

• First ionization energy

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Periodic Table of Video

• http://www.periodicvideos.com/