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chapter 25-periodic properties Flashcards

1
Q

what is the periodic law

A

which states that the chemical properties of the elements are dependent in a systematic way upon their atomic numbers

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2
Q

what does a group represent

A

elements that have the same electronic configuration in their valence or outermost shell

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3
Q

what are valence electrons

A

electrons in the valence shell

-these electrons are involved in chemical bonding and determining the chemical reactivity and properties of the element

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4
Q

what are A elements

A

representative element

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5
Q

what are representative elements

A

which have either s or p sublevels as their outermost orbitals
-groups IA through VIIA all of which have incompletely filled s or p subshells of the highest principal number

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6
Q

what are B elements

A

nonrepresentative elements

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7
Q

what are non representative elements

A

include the transition elements which have partly filled d sublevels and the lanthanide and actinide series which have partly filled f subshells

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8
Q

all elements seek to gain or lose what and why

A

valence electrons so as to achieve the stable fully-filled formations possessed by the inert or noble gasses of Group VIIIA

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9
Q

what are some periodic trends

A
  • left to right across a period: protons are added one at a time and the electrons of the outermost shell experience an increasing degree or nuclear attraction becoming closer and more tightly bound to the nucleus
  • top to bottom down a given column the outermost electrons become less tightly bound to the nucleus
  • both these trend show that Zeff is at a max for elements in the top-right of the table and at a min for those in the bottom-left and help explain elemental properties such as atomic radius, ionization potential, electron affinity, and electronegativity
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10
Q

why going top to bottom are the e become less tightly bound to the nucleus

A

bc the number of filled principal energy levels (which shield the outermost electrons from attraction by the nucleus) increases downward within each group

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11
Q

what is effective nuclear charge (Zeff):

A

net positive charge from the nucleus as felt by an electron

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12
Q

what is the atomic radius of an element equal

A

one-half the distance between the centers of 2 atoms of that element that are just barely touching each other

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13
Q

atomic radii trend

A

decreases across a period from L to R and increases down a given group

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14
Q

what occupies the majority of the volume of an atom

A

the electron cloud

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15
Q

will changing the size of the nucleus affect the size of the atom yes or no and why

A

no the size of the electron cloud will change the radius of an atom but changing the size of the nucleus will not directly affect the size of the atom

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16
Q

atomic size trend in a period and the reason

A

Left to Right

  • electrons are added 1 at a tie to the outer energy shells
  • e within the same shell do not shield one another from the attractive pull of protons so since the number of protons is also increasing from L to R across a period the Zeff inceases as well
  • the greater the + charge experienced by the valence e the greater the Zeff the closer those e are pulled towards the nucleus and the smaller the atomic radius
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17
Q

atomic size trend in a group and the reason

A

top to bottom

  • the number of e and filled e shells increases but the # of valence e within a group remain the same the valence e will be found farther from the nucleus as they are in progressively larger energy shells
  • Zeff will become smaller with distance so valence electrons in higher energy shells will feel less pull from the nucleus
  • With more e comes increased repulsion from the additional negative charges therefore increasing the atomic radius
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18
Q

what is ionic radius

A

radius of a cation or anion

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19
Q

are cations smaller or larger than the corresponding neural atoms and why

A

small than corresponding neutral atoms since fewer e leads to less repulsion amount the remaining e

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20
Q

are anions smaller or larger than the corresponding neural atoms and why

A

larger than corresponding neutral atoms bc having a greater number of e causes more repulsion resulting in a greater distance between electrons and a larger atomic radius

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21
Q

what is ionization energy

A

is the energy required to completely remove an electron from a gaseous atom or ion

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22
Q

trends for ionization energy

A
  • The closer and more tightly an e is to the nucleus the more difficult it will be to remove and the higher the ionization energy
  • Increases left to right across a period as Zeff increases
  • Moving down a group the ionization energy decreases as Zeff decreases
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23
Q

what is First ionization energy

A

is the energy needed to remove a second valence electron from the univalent ion to form the divalent ion

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24
Q

what is Second ionization energy

A

the energy needed to remove a second valence electron from the univalent ion to form the divalent ion

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25
what is the trend from group IA ionization elements and why
low ionization energy bc the loss of an electron results in the formation of a stable noble-gas configuration
26
what is electron affinity (EA)
The energy change that occurs when an electron is added to a gaseous atom and it represents the ease with which the atom can accept an electron
27
o A positive electron affinity value represents energy release when an electron is added to an atom ( more common used) OR A negative electron affinity represents a release of energy
both are true
28
EA for Group IA or alkaline earth metals
-have low electron affinity | oTheir elements are relatively stable bc their s subshell is filled
29
EA for group VIIA or halogens
have high electron affinities bc the addition of an e to the atom results in a completely filled shell which represents a stable e configuration
30
EA for group VIIIA or noble gases
have electron affinities on the order of 0 bc they already possess full shells and cannot readily accept electrons
31
what is electronegativity
a measure of the attraction an atom has for electrons in a chemical bond
32
which element is the most electronegativity
F
33
what is the electronegativity trend
- Related to Zeff low Zeff will have low electronegativities bc their nuclei don’t attract electrons strongly and high Zeff will have high electronegativities bc of the strong pull the nucleus has on electrons - Increase L to R across periods and decreased Top to Bottom down a group
34
what are metals
shiny solids ( exept mercury) at room temperature and generally have high melting points and densities
35
characteristics of metals
-Can deform without breaking Characteristicss: - Large atomic radius, low ionization energy, low electronegativity all due to the fact that the few e in the valence shell of a metal atom can be removed -Good conductors of heaty and electricity bc valence e can move freely
36
malleability
metals can be hammered into shape
37
Ductility
ability to be drawn into wires
38
what are non-metals
generally brittle in the solid state and show little no metallic luster
39
what are non-metals characteristics
- High ionization energies and electronegativities - Poor conductors of heat and electricity - They share the ability to gain e easily but have a wide range of chemical behaviors and reactivates
40
what are the metalloids characteristics
- Densities, bp, and melting points fluctuate widely - Electronegativities and ionization energies of metalloids lie between those of metals and non metals therefore these elements possess characteristics of both those classes
41
what are Polyatomic anions
contain O and therefore are called oxyanions
42
naming of polyatomic
-The one with the less O ends in ite and the one with more is ate o NO2- nitrite o NO3nitrate
43
mental of L side form what type of ions
positive
44
non-metals on R side of table form what type of ion
negative
45
what is the charge of the groups that form monatomic ions
``` oAlkali metals (Group IA) usually form cations with a single positive charge oalkaline earth metals ( Group IIA) form cations with a double positive charge oHalogens ( Group VIIA) form anions (halides) with a single negative charge ```
46
what group are alkali metals
Group IA
47
properties of alkali metals
possess most of the physical properties common to metals yet their densities are lower than those of other metals - have only 1 loosely bound e in their outermost shell giving them the largest atomic radii of all the elements in their period - their metallic properties and high reactivity are due to their low ionization energies they lose their valence e to form univalent cations allow them to easily form +1 cations - low electronegativities and react very readily with nonmentals especially halogens
48
group are the alkaline earth metals in
group IIA
49
what are the properties of alkaline earth metals
- posses many characteristically metallic properties just like alali these properties are dependent on the ease with which they lose e - have 2 e in their outer shell and have smaller atomic radii than the alkali metals but the 2 valence 2 e are not held very tightly by the nucleus so can be removed to form divalent cations so have a +2 charge - low electronegativities and positive electron affinities
50
what group is C in
Group IVA
51
properties of carbon
that explains why C tends not to form ions which would need to be +4 or -4 to reach noble gas but rather participates in e sharing -Is usually most stable with 4 covalent bods
52
what is Pnictogens group
Group VA
53
properties of N
- often forms covalent bonds but most commonly forms 3 per atom Holds a + charge in organic rxn making several N containing cpds good bases
54
what group are Chalcogens in
Groups VIA contains O
55
characteristics of Chalcogens
Group VIA which contains O is characterized by elements requiring 2 additional valence electrons to complete their outermonst shells - They tend to fairly electronegative and to form -2 anions but they can also participate in covalent bonds preferring to have 2 shared electron paris and 2 nonbonded pairs
56
what group are Halogens
VIIA
57
characteristics of Halogens
-Highly reactive non-metals with 1 valence electron less than the closest noble gas -Form -1 anions -Highly variable in their physical properties oCan be gases ( F2 and Cl2) liquid (B2) and solid (I2) at room temperature -Chemical properties are more uniform oHigh electronegativites oReactive towards alkali metals and alkaline earth which want to donate electrons to halogens to form stable ionic crystals
58
what group are noble gases in (inert gas)
group VIIIA
59
characteristics of noble gases
- Unreactive bc they have a complete valence shell which is an energetically favored arrangement - High ionization energies - No electronegativities - Low bp - Gasses at room temperature
60
what group are transition elements in
Group IB to VIIB
61
characteristics of transition elements
- All metals - Hard and have a high melting point and boiling point - L to R the 5d orbital becomes progressively more filled - chemically they have low ionization energy and may exist in a variety of positively charged forms (oxidation state)
62
bc of what allows transition elements to form many different ionic and partially ionic cpds
their ability to attain positive oxidation states
63
formation of what causes the d orbital to split into 2 sub levels
complexes - enables many of the complexes to absorb certain frequencies of light and the frequencies not absorbed are called suntraction frequencies which give complexes their characteristic colors

DAT-chemistry (4 decks)

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