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Chemistry > Unit 3 > Flashcards

Unit 3 Flashcards

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1
Q

Valence Electrons

A

Electrons that are found in an atom’s outermost energy level (furthest from the nucleus) these are important because they determine how an atom will bond and react

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2
Q

The Octet Rule

A

All elements are trying to become stable by achieving a full valence shell of electrons

  • Each elecment wants to be as stable as possible by obtaining a full valence shell
  • Most of the elements want to have 8 valence electrons
  • Atoms will form an ion (atom with a charge) based on the easiest way to obtain a full shell
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3
Q

Group

A

Also called a family
VERTICAL columns on the periodic table
Elements in the same group have the same number of valence electrons and similar properties

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4
Q

Period

A

HORIZONTAL rows of the periodic table

Have the same number of energy levels

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5
Q

What are the 3 main element classifications?

A

Metal
Non-metal
Metalliods

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6
Q

Metals

A 
Examples: Sodium, Calcium, Gold
Common properties:
-Shiny
-Usually grey
-High melting/boiling points, mostly solid at room temp.
-Good conductors of heat and electricity
-Reacts with acid to form hydrogen gas
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7
Q

Nonmetals

A

Examples: Argon, Chlorine, Neon
Common properties:
-Mostly dull looking
-Most are solid, but some liquids and even gases at room temp
-Not good conductors of heat and electricity
-Low melting/boiling points

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8
Q

Metalloids

A

Examples: Boron, Silicon, Antinomy
Common properties:
Have some properties in common with metals and some with nonmetals

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9
Q

Group 1

A

Alklai Metals
Lithium, Sodium, Potassium, Rubidium, Cesium, Francium
-Does NOT include hydrogen
-soft metals-usually exist in compounds
-highly reactive-react voilently with water
valence electrons: 1
Ion formed: +1

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10
Q

Group 2

A

Alklai Earth Metals
Beryllium, Magnesium, Calcium, Strontium, Barium, Radium
-reactive metals (slightly less reactive than Group 1 since they have to lose 2 valence electrons)
-Minerals-important in the body
valence electrons: 2
Ions formed: +2

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11
Q

Group 13

A

Boron family
Examples: Boron, Aluminum, Gallium, Indium, Thallium
-aluminum is the most abundant metal and will form a +3 ion
valence electrons: 3
Ions formed: +3

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12
Q

Group 14

A 
Carbon family
Carbon, Silicon, Germanium, Tin, Lead
-Carbon is the base for life (organic compounds)
-Silicon is used in computer chips
number of valence electrons: 4
Ion formed: +/- 4
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13
Q

Group 15

A

Nitrogen family
Nitrogen, Phosphorus, Arsenic, Antimony, Bismuth
-2 nonmetals (Nitrogen and Phophorus) 2 metalloids (Arsenic and Antimony) and 1 metal (Bismuth)
Number of valence electrons: 5
Ions formed: -3 ion

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14
Q

Group 16

A

Oxygen family
Oxygen, Sulfur, Selenium, Tellurium, Polonium
-3 nonmetals (Oxygen, Sulfur, Selenium), 1 metalloid (Tellurium) and and 1 metal (Polonium)
Number of valence electrons: 6
Ions formed: -2

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15
Q

Group 17

A 
Halogens
Flourine, Chlorine, Bromine, Iodine, Astatine, 
-Most reactive group of nonmetals
-Found in nature as diatomic element
 I Bring Clay From Our New House
-Readily combine with metals to form salt (NaCl)
Number of valence electrons: 7
Ion formed: -1
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16
Q

Group 18

A 
Noble Gases
Helium, Neon, Krypton, Xenon
-gases
-non reactive (8 valence electrons)
-Helium has 2 valence electrons on outershell
Number of valence electrons: 8
Ions formed: X
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17
Q

Group 3-12

A

Transition Metals

  • When they form ions, they become more complicated because they have multiple energy levels involved, therefore produce variable charges
  • Also includes metals under the staircase
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18
Q

Two groups at the bottom of the periodic table are known as ____

A

Rare Earth Metals

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19
Q

Period trends

A

The pattern for a horizontal row on the periodic table (left to right)

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20
Q

Group trend

A

The pattern for a vertical column on the periodic table (up and down)

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21
Q

Atomic radius

A

The measure of the size of an atom

measured from the nucleus to the outermost electrons

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22
Q

Period trend for atomic radius

A

As you move across a period (left to right) atomic radius decreases

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23
Q

Group trend for atomic radius

A

As you move down a group, the atomic radius increases

24
Q

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

A

As you move towards the right, the nuclear charge increases. The greater the nuclear charge, the more tightly a nucleus is holding onto the valence shell. Being more tightly packed means they are smaller

25
Q

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

A

Each time you move down a group, the atomic radius gets bigger because you add energy levels, therefore the outermost rings are far from the nucleus

26
Q

Ion radius

A

The size of an atom

27
Q

Positive Ion

A

Cations

When atoms lose electrons, they get smaller

28
Q

Why do atoms get smaller when they lose electrons?

A

They lose their valence electrons, which means they have fewer energy levels

29
Q

Negative Ions

A

Anions

When an atom gains electrons, it gets bigger

30
Q

Why do atoms get bigger when they gain electrons?

A

When an atom gains electrons, it holds onto more electrons less tightly

31
Q

Ionization energy

A

The energy required to remove an electrons from an atom

32
Q

Period trend for ionization energy

A

As you move across a period, ionization energy increases

33
Q

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

A

As you move across a period the nuclear charge increases. The greater the nuclear charge, the more tightly the atom holds onto the valence electrons. The greater the nuclear charge, the more energy it will take to remove an electron

34
Q

Group trend for ionization energy

A

Ionization energy decreases as you move down a group

35
Q

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

A

As you move down a group, the atom gets bigger. The number of energy levels increases and the valence shell begins to get further away from the nucleus. The farther away the rings are from the nucleus, the less charge it will take to remove the electron

36
Q

Orbital

A

3D representation of the volume where an electron can be found

37
Q

How are electrons represented in orbitals?

How many electrons can each individual orbital hold?

A
  • By arrows

- each can hold 2

38
Q

What are the four main shapes?

A

s, d, p, and f

39
Q

s

A

1 orbital

Can hold 2 electrons

40
Q

p

A

3 orbitals

Can hold 6 electrons

41
Q

d

A

5 orbitals

Can hold 10 electrons

42
Q

f

A

7 orbitals

Can hold 14 electrons

43
Q

Shapes on energy level 1

A

s

44
Q

Shapes on energy level 2

A

s, p

45
Q

Shapes on energy level 3

A

s, p, d

46
Q

Shapes on energy level 4

A

s, p, d, f

47
Q

Shapes on energy level 5

A

s, p, d, f

48
Q

Shapes on energy level 6

A

s, d, p

49
Q

Shapes on energy level 7

A

s, p

50
Q

Aufbau principle

A

Electrons will always fill the lowest energy level first

Lazy tenement

51
Q

Pauli Exclusion Principle

A

No 2 electrons can occupy an orbital unless they spin in opposite directions
Shoebox rule

52
Q

Hund’s rule

A

Electrons spread out within an orbital until they must pair up
City bus

53
Q

Ground state

A

The lowest energy state

54
Q

Excited state

A

A state in which an atom has more energy than it does in the ground state

55
Q

Electron configuration

A

Arrangement of electrons in atoms

Chemistry (11 decks)

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