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IB Chemistry > Unit 3: Periodicity > Flashcards

Unit 3: Periodicity Flashcards

1
Q

Properties of Metal

A
  • usually a solid @ room temp (except Hg)
  • shiny/lustrous
  • malleable: can be hammered into sheets
  • good conductors of heat + electricity
  • high m.p.
  • ductile: can be drawn into wires
  • lose e-
  • low values for electronegativity, ionization energy, and electron affinity values
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2
Q

Properties of Nonmetals

A
  • dull/lack luster
  • poor conductor
    -insulator of heat + electricity
  • gain e-
  • high values for electronegativity, ionization energy, and electron affinity values
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3
Q

Properties of Metalloids

A
  • Have characteristics of both metals & nonmetals
  • Physical properties/ appearance resemble metals
  • Chemical properties resemble nonmetals
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4
Q

Periods

A

Horizontal Rows (1-7)

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5
Q

Groups

A

Vertical columns (1-18)
- sometime groups 3-12 are “omitted” and we call the remaining groups, Groups 1A-8A

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6
Q

Group 1

A

Alkali metals

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7
Q

Groups 3-12

A

Transition metals (d-block)

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8
Q

Group 17

A

Halogens

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9
Q

Group 18

A

Noble gasses

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10
Q

Elements 58-71

A

Lanthanoids (f-block)

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11
Q

Elements 90-103

A

Actinoids (f-block)

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12
Q

Atomic Radius (group-trend)

A

As you move down a group, atomic radius increases
(“snowman effect”

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13
Q

Why does atomic radius increase as you move down a group?

A

As the # of occupied energy levels increases, so does the size of the atom

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14
Q

Atomic Radius (period-trend)

A

As you move across a period, atomic radius decreases

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15
Q

Why does the atomic radius decrease as you move across a period?

A

Same: e- are being added to the same energy level
Different: the # of protons increases so the attraction between the increasingly positive center and (-) e- increases. This attraction “pulls in” the outer e-

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16
Q

Ionic Radius (Cations) (Group)

A

Cations increase in size down a group as the outer energy level gets further from the nucleus (less attraction)

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17
Q

Ionic Radius (Cations) (Period)

A

Cations in a period contain the same # of e- but an increasing # of protons, creasing the attraction between the (+) center and the (-) e-. thus decreasing the radius

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18
Q

Ionic Radius (Anions) (Group)

A

Anions increase in size down a group as the outer energy level gets further from the nucleus (less attraction)

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19
Q

Ionic Radius (Anions) (Period)

A

Anions in a period contain the same # of e- but an increasing # of protons, increasing the attraction between the (=) center and the (-) e-, thus decreasing the radis

20
Q

First Ionization Energy (definition)

A

The energy required to remove one electron from an atom in its gaseous state

21
Q

Ionization Energy (Group) (Trend)

A

As you move down a group, ionization energy decreases

22
Q

Why does the first ionization energy decrease as you move down a group?

A

The # of occupied energy levels increases, so the increased distance between the e- and the nucleus reduces the attraction (lower I.E. makes it easier to remove e-)

23
Q

Ionization Energy (Period) (Trend)

A

As you move across a period, ionization energy increases

24
Q

Why does the first ionization energy increase as you move across a period?

A

The extra e- are filling the same energy level, but the extra protons in the nucleus attract the energy level close making it harder to remove an e- (increasing I.E.)

25
Q

Electronegativity (definition)

A

A measure the ability to attract electrons in a covalent bond

26
Q

Electronegativity (Group) (Trend)

A

As you move down a group, electronegativity decreases

27
Q

Why does electronegativity decrease as you move down a group?

A

The bonding e- (outermost e-) are getting father from the nucleus and so there is reduced attraction

28
Q

Electronegativity (Period) (Trend)

A

As you move across a period electronegativity increases

29
Q

Why does electronegativity increase as you move across a period?

A

There is an increase in # of protons, resulting in an increased attraction between the nucleus and the bounding e-

30
Q

Electronegativity (Noble gasses)

A

no values
no desire to attract an e-

31
Q

Electron Affinity (definition)

A

The energy change when an electron is added to an isolated atom in the gaseous state

32
Q

Metalloids

A

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

33
Q

Melting Point

A

depends on the structure of the element and the type of attractive forces which hold the atoms together

34
Q

Melting Point (Metals)

A

exhibit metallic bounding (ex. Na, Mg, Al) which increase in strength as the # of valence e- increases

35
Q

Melting Point (Silicon)

A

The metalloid has a “unique” covalent structure with very strong bonds resulting in high m.p.

36
Q

Melting Point (Nonmetals)

A

P, S, Cl have a simpler molecuar structures with weak forces of attractions

37
Q

Melting Point (Ar/noble gasses)

A

exist as monatomic molecules (single atoms) with extremely weak forces of attraction between the atoms

38
Q

Melting Point (Group 1)

A

m.p. decreases down the group as the atoms become larger and the strength of the metallic bond decreases

39
Q

Melting Point (Group 17)

A

Halogens are held together by London dispersion forces. These get stronger as the number of electrons in the molecule increases

40
Q

Alkali metals

A
  • very reactive metals
  • readily lose e- (low I.E.)
    -Chemical reactivity increases as you go down the group since less E is required to remove the e-, aiding in a faster/more vigorous reaction
41
Q

Halogens

A
  • very reactive non-metals
  • Exist as diatomic molecules
  • react by gaining an e-
  • Reactivity decreases as you go down the group as the outer shell gets further from the nucleus, decreasing the attraction for another e-
42
Q

Most reactive metal

A

Francium

43
Q

Most reactive nonmetal

A

Fluorine

44
Q

Acid

A
  • starts w/ H
  • lots of H+
    (0-7)
45
Q

Basic

A
  • end w/ OH
  • lots of OH- (hydroxide)
    (pH>7)
46
Q

Period 3

A

basic-> amphoteric -> acidic

IB Chemistry (2 decks)

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