(Cb) Chapter 2: The Periodic Table

Front 1

Periodic Law

Back 1

The notion that if the elements are put in an order of atomic number, they will exhibit trend in their physical and chemical properties

Front 2

Periods

Back 2

Rows. They go across the periodic table. One way to remember this is to imagine that row as sentences and periods at the end of these sentences

Front 3

Groups/Families

Back 3

Columns. They go from top to bottom

Front 4

Which electrons have the highest potential energy?

Back 4

The valence shell electrons. they determine an elements reactivity

Front 5

What do the seven periods correspond to?

Back 5

They correspond to seven principle quantum numbers

Front 6

Representative Elements

Back 6

These elements are the designated with the letter A on the periodic table. They refer to elements that fill the s and p orbitals

Front 7

Non representative elements

Back 7

These elements are designated with the letter B on the periodic table. They refer to the transition elements the fill the s p and f orbitals. They include the lanthanides and actinides

Front 8

Z Effective Nuclear Charge

Back 8

This is the attraction between the outer shell electrons and the positive charge of the nucleus

Front 9

Atomic size across the periodic table

Back 9

The size of the atom will get smaller as you move across the periodic table. The reason for this is that the effective nuclear charge gets stronger with each additional element you are adding more and more protons.

Front 10

Atomic size down the periodic table

Back 10

The size of the atom will increase. The reason for this is that the effective nuclear charge will get weaker as you move down the table given that you are adding quantum number. The additional orbitals means that the electrons will be further away from the nucleus.

Front 11

How do you measure the size of an atom?

Back 11

You take one half the distance between the nuclei of two atoms of the same element.

Front 12

Ionization Energy across the periodic table

Back 12

The ionization energy will increase as you move across the periodic table. The reason for this is that as you move across the table the effective nuclear charge of the atom will increase. This means that the hold on the valence electrons will be stronger and thus raise the ionization energy.

Front 13

Ionization energy

Back 13

The energy needed to remove an electron from an atom in the gas phase. The higher the ionization energy the more energy is needed to remove an electron

Front 14

Ionization energy down the periodic table

Back 14

The ionization energy will decrease as you move down the periodic table. The reason for this that the effective nuclear charge of the atom decreases since the there are more principle quantum numbers that increase the distance between outer shell electrons that the nucleus

Front 15

Natural State of group 1 and 2 elements

Back 15

The natural state of group one and group two element is that they usually found ionically bonded to other elements. As in that are usually found with other elements and never on their own.

Front 16

What type of reaction is it when an atom loses an electron and why?

Back 16

It is an endothermic reaction because you need to invest energy energy into the atom to remove the electron. This means that energy is going into the atom making the reaction endothermic

Front 17

What type of reaction is it when an atom gains an electron and why?

Back 17

It is an exothermic reaction. The reason why is that as the atom gains the electron, energy is released from atom making it an exothermic reaction.

Front 18

Electron affinity across the periodic table

Back 18

As you move across the table electron affinity increases. The reason being the increase in effective nuclear charge as you move across the table

Front 19

Electron affinity down the periodic table

Back 19

As you move down the table electron affinity decreases given the decrease in the effective nuclear charge

Front 20

Electronegativity

Back 20

This is the attractive force that an atom will exerts on electrons IN A BOND. This is an important distinction. It not how much an atom will attract and electron but how much it will attract and electron in a bond.

Front 21

Electron Affinity

Back 21

Electron affinity is not just how much an atom wants an electron. Electron affinity is how much energy is released as an atom gains an electron.

Front 22

Electronegativity across the periodic table

Back 22

Electronegativity will increase as you move across the table. F being the most electronegative atom

Front 23

Electronegativity down the periodic table

Back 23

Electronegativity will decrease as you move down the table.

Front 24

Malleability

Back 24

The ability of a metal to be hammered into different shapes

Front 25

Ductility

Back 25

The ability of a metal to be draw into a wire

Front 26

Melting points of metal

Back 26

Metals have higher melting points

Front 27

Conductive ability of metals

Back 27

Metals have very good conductive abilities

Front 28

Oxidation states of metal

Back 28

Some metals can have oxidation states of +1 or +2 since metals very readily give up their electrons

Front 29

Density of metals

Back 29

Metals generally have high densities. One exception is lithium.

Front 30

Which metal is liquid under standard conditions

Back 30

Mercury

Front 31

Texture of non metals

Back 31

Non metals are brittle in the solid state and lack metallic luster

Front 32

Non metal conduction properties

Back 32

Non metals do not conduct heat or electricity well.

Front 33

Metalloids

Back 33

These elements have properties of both metals and non metals

Front 34

Alkali Metals

Back 34

They form +1 cations since they easily give up one electron. They have lower densities and react readily with non metals

Front 35

Alkaline Earth Metals

Back 35

They for +2 cations since they easily give up two electrons.

Front 36

Active metals

Back 36

Alkali and Alkakine Metals are both very reactive with non metals which is why they are called active metals

Front 37

Halogens

Back 37

Very reactive non metals. They react very readily with the group one and two metals. Their physical properties range from solid to gas at room temperature

Front 38

Why are halogens so reactive?

Back 38

Halogens are so reactive because they have seven valence electrons and are 'looking' to find one more electron to achieve the noble gas configuration

Front 39

Noble gases

Back 39

Very unreactive

Front 40

Oxidation states of transition metals

Back 40

given the fact that transition metals can form a number of different oxidation states they can form a number of different type of bonds. they can form water complexes and different ionic bond.