Chapter 25 - Periodic Properties Flashcards Preview

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Flashcards in Chapter 25 - Periodic Properties Deck (27):
1

Who published the first periodic table? What did it show

Dmitri Mendeleev - showed the ordering of elements according to atomic weight producing a pattern where similar properties periodically recurred

2

Who reorganized the period table? Based on what?

Henry Moosely - based on increasing atomic number

3

What is the periodic law?

States that the chemical properties of the elements are dependent, in a systemic way, upon their atomic numbers

4

How is there periodic table organized today?
- What do they represent?
- What does the Roman numerals represent?

In periods (rows) and groups (columns)
- There are 7 periods that represent the principal quantum numbers n = 1 to n = 7 and each period is filled sequentially; groups represent elements that have the electronic configuration in valence shell and share similar chemical properties
- They are located above each group and represent the number of valence electrons

5

What are representative and non-representative/transition elements?

1) Representative: Have s or p sub-levels as their outermost orbitals - Groups IA through VIIA which have incompletely filled s or p sub-shells

2) Non-representative/Transition: Have partly filled d sub-levels, and the lanthanide and actinide series, which have partly filled f sub-levels

6

What do all elements seek to do?

Gain or lose electrons to achieve a fully filled formation possessed by the inert or noble gases of Group VIIIA

7

What are the two trends exhibited in the periodic table?
* What does Zeff show?

1) From left to right across a period, protons are added one at a time and the electrons of the outermost shell experience an increasing degree of nuclear attraction, becoming closer and more tightly bound to the nucleus
- Net positive charge from the nucleus is called the effective nuclear charge (Zeff)

2) From top to bottom down a given column, the outermost electrons become less tightly bound to the nucleus because the number of filled principle energy levels increases downward within each group

*Together, the trends show that Zeff is at a maximum for elements top-right and at a minimum for those in the bottom-left

8

What is the atomic radius equal to?

One-half the distance between the centres of two atoms of that element that are just barely touching each other

9

What is the trend for atomic radius?

Atomic radius decreases across a period from left to right and increases down a given group
- Largest = bottom-left

10

What happens to an atom from left to right in regard to atomic radii?

Electrons are added left to right one at a time to the outer energy shell
- Electrons within the same shell do not shield one another from the attractive pull of protons, therefore since the number of protons are increasing left to right, the effective nuclear charge of Zeff increases as well
- The greater positive charge experienced by the valence electrons, the closer those electrons are pulled toward the nucleus and the smaller the atomic radius

11

What happens from top to bottom in regard to atomic radii?

The number of electrons and filled electron shells increase from top to bottom in a group
- Although the number of valence electrons within a group remain the same, these valence electrons are found further and further from the nucleus as they are in larger energy shells
- Zeff becomes smaller with distance, so valence electrons in higher energy shells will feel less pull from the nucleus
- More electrons comes increased repulsion from the additional negative charge
- Atomic radii increases

12

What is ionic radius?

The radius of a cation or anion
- Cations = smaller since they possess fewer electrons creating less repulsion
- Anions = larger since they have greater number of electrons creating greater repulsion = larger distance between electrons and larger atomic radius

13

What is ionization energy?
- Difficulty?

Energy required to completely remove an electron from a gaseous atom or ion which requires an input of energy (endothermic)
- The closer and more tightly bound an electron is to the nucleus, the more difficult it is to remove and the higher the ionization energy

14

What is the first ionization energy? Second?

Energy required to remove one valence electron from the parent atom
- It is the energy needed to remove a second valence electron from the univalent ion to form the divalent ion, and so on

15

Is the second ionization energy always greater than the first? Why?

Yes - they grow increasing larger with each ionization energy

16

What is the ionization energy trend?

Increases left to right across as period as Zeff increases

Moving down a group, decreases as Zeff decreases

17

Why does Group IA have low ionization energy?

Because the loss of an electron results in the formation of a stable, noble-gas configuration

18

What is electron affinity?
- What is Zeff effect on affinity?

The energy change occurs when an electron is added to a gaseous atom, and it represents the ease with which the atom can accept and electron
- The higher the Zeff, the greater the electron affinity will be

19

There are two sign conventions used in discussing electron affinity, which is the more common?

States that a positive electron affinity value represents energy release when an electron is added to an atom

Ex. X(g) + e- = X- (g)

20

What is the less common sign convention in regard to electron affinity?

A negative electron affinity represents a released of energy

21

What are some generalizations about electron affinity?

- Group IIA (alkaline earth metals) have low electron affinity values because they are relatively stable due to their s sub-shell being filled
- Group VIIA (halogens) have high electron affinity because the addition of an electron results in a completely filled shell; a stable configuration
- Group VIIIA (noble gases) have electron affinity of zero because they possess a full shell and cannot accept electrons

22

What is electronegativity?
- Trend?

A measure of the attraction an atom has for electrons in a chemical bond - the greater it is, the greater the attraction for bonding electrons
- Trend: Increases from left to right across periods and decreases from top to bottom down a group

23

What is the Pauling electronegativity scale?
- Low Zeff?
- High Zeff?

Ranges from 0.7 (most electropositive - Ce) to 4.0 (more electronegative - F)
- Low Zeff = low electronegativity because nuclei does not attract electrons strongly
- High Zeff = high electronegativity because of the strong pull the nucleus has on electrons

24

List the types of elements

1) Metals
2) Non-metals
3) Mettaloids

25

Review the nomenclature of cations and anions

Page 281/282

26

Describe ion charges of metals and non-metals

Metals: Generally form positive ions

Non-metals: Generally form negative ions

27

List the groups of the periodic table
- What does have a group?

1) Alkali Metals (Group IA)
- 1 loosely bound electron in their outermost shell
- Largest atomic radii
- Highly reactive due to their low ionization energies
- Form +1 cations

2) Alkaline Earth Metals (Group IIA)
- 2 electrons in their outermost shell
- Form divalent cations (+2 cations)

3) Carbon Group (Group IVA)
- Wide range of characteristics and includes metals, non-metals, and metalloids
- 2 electrons in outermost p orbital
- Participates more so in electron sharing; most stable with 4 covalent bonds

4) Pnictogens (Group VA)
- Wide variety of characteristics and includes non-metals, metals, and metalloids
- Forms best with 3 covalent bonds

5) Chalcogens (Group VIA)
- Contains oxygen and is characterized by elements requiring 2 additional valence electrons to complete their outermost shell
- Fairly electronegative and forms -2 anions
- May participate in covalent bonding, preferring 2 shared electron pairs and 2 non-bonded pairs

6) Halogens (Group VIIA)
- Highly reactive non-metals with 1 valence electron less than closest noble gas
- Range from (g) to (l) to (s)
- High electronegativity (highest is F)

7) Noble Gases (Group VIIIA)
- Most stable because have complete valence shell
- High ionizagtion and no electronegativity
- Low boiling point and gas at room temperature

* Transition Elements (Group IB to VIIIB)
- Considered mental
- Hard, high melting/boiling points
- Loose electron in outer orbital contributes to malleability
- Positively charged or oxidative states