Ch. 2: The Periodic Table (Complete) Flashcards
CH 2.3
what are the three key rules that control how valence electrons work in an atom?
- L to R across a period = electrons and protons are added one at a time
- Down the elements of a group = principle quantum number increases by one each time
- elements can gain or lose electrons in order to achieve a stable octet formation representative of noble (inert) gases
explain why: electrons and protons are added one at a time L to R across a period and what the effect is
- as the positivity of the nucleus increases, the electrons surrounding the nucleus experience a stronger electrostatic pull toward the atom’s center
- this causes the electron cloud (the outer boundary defined by the valence shell electrons) moves closer and binds more tightly to the nucleus
defn: effective nuclear charge (Zeff)
what is the effect of nonvalence electrons on Zeff
a measure of the net positive charge experienced by the outermost electrons in the electrostatic attraction between the valence shell electrons and the nucleus
the pull is somewhat mitigated by nonvalence electrons residing closer to the nucleus
how does Zeff change from L to R for elements in the same period?
Zeff increases
what are the implications of the principal quantum number increasing down the elements of a given group? (5)
- the valence electrons are increasingly separated from the nucleus by a greater number of filled principal energy levels (inner shells)
- this means a reduction in the electrostatic attraction between the valence electrons and the positively charged nucleus
- the outermost electrons are held less tightly
- increased shielding created by inner shell electrons cancels the increased positivity of the nucleus
- thus Zeff is pretty constant among the elements within a group
defn: octet rule
elements (esp. the ones with biological roles) tend to be most stable with 8 electrons in their valence shell, although there are many exceptions to this rule
basis + defn: atomic radius
basis: imagine an atom as a cloud of electrons surrounding a dense core of protons and neutrons
atomic radius = one-half of the distance between the centers of two atoms of an element that are briefly in contact with each other
what happens to atomic radius L to R across a period and why (4)?
atomic radius DECREASES L to R across a period
- protons and electrons are added one at a time to the atoms
- electrons are only added to the outermost shell, the number of inner-shell electrons remains constant
- this increases the positive charge of the nucleus
- which then pulls the outer electrons more closely inward and holds them more tightly
what happens to atomic radius down a group and why (4)?
atomic radius INCREASES down a group
- increasing principal quantum number down a group –> 2. valence electrons will be found further from nucleus 3. bc the # of inner shell electrons is increasing –> 4. separating the valence shell from nucleus
what two generalizations must be made to understand ionic radii?
- metals lose electrons and become positive; nonmetals gain electrons and become negative
- metalloids can go in either direction, but tend to follow the trend based on which side of the metalloid line they fall on
what is the trend (2) in ionic radii for nonmetals close to the metalloid line in comparison to nonmetals close to group VIIIA?
nonmetals close to the metalloid line:
1. require more electrons than other nonmetals to achieve octet (gain electrons)
2. have LARGER ionic radius than counterparts
what is the trend in ionic radii for metals close to the metalloid line in comparison to metals close to group VIIIA?
metals close to the metalloid line:
1. have more electrons to lose to achieve octet (lose electrons)
2. have much SMALLER ionic radii than other metals
defn and aka: ionization energy
aka: ionization potential
the energy required to remove an electron from a gaseous species
explain the process associated with “ionization energy”. is this an endothermic or exothermic process? why?
ionization energy = energy required to remove an electron from a gaseous species (SO PROCESS = IONIZATION)
this is an ENDOTHERMIC process because removing an electron from an atom always requires heat input
how does Zeff impact ionization energy?
what does this mean in terms of the trend of ionization energy on the periodic table?
- the greater the atom’s Zeff (the closer the valence e’s are to the nucleus) –> 2. the more bound they are –> 3. the more difficult to remove one or more electrons –> 4. increases ionization energy
Ionization energy INCREASES from L to R across a period and from bottom to top in a group
expln + defn: first ionization energy and second ionization energy
the subsequent removal of a second or third electron requires increasing amounts of energy because the removal of more than one electron means that electrons are being removed from an increasingly cationic (positive species)
first ionization energy = the energy necessary to remove the first electron
second ionization energy = the energy necessary to remove the second electron from the univalent cation (X+) to form the divalent cation (X2+)
which is larger: the first or second ionization energy?
the second ionization energy
groups + natural state + expln for name: active metals
elements in groups 1 and 2
called active metals bc they have such low ionization energies
do not naturally exist in neutral forms, but are always found in ionic compounds, minerals, or ores
what changes must be made to group IA (alkali metals) or group IIA (alkaline earth metals) to form a stable, filled valence shell?
group IA –> the loss of one electron
group IIA –> the loss of 2 electrons
why are the second ionization energies for group 1 monovalent cations disproportionally larger than the ones for group 2 or other monovalent cations?
removing one electron from a group 1 metal results in a noble gas-like electron configuration
why are the ionization energies of nobel gases so high?
they are the least likely to give up electrons, have a stable electron configuration already
what is the periodic trend of ionization energy?
L to R increases
bottom to top INCREASES
defn: electron affinity
the energy dissipated by a gaseous species when it gains an electron
the opposite concept of ionization energy
is the process associated with electron affinity endothermic or exothermic?
exothermic –> gaining an electron expels energy in the form of heat
deltaHrxn has a negative sign, but electron affinity is reported as positive bc it refers to the energy dissipated
why are halogens the most greedy elemental group when it comes to electrons?
they only require one additional electron to complete their octet
what is the periodic trend for electron affinity? why (@)?
L to R = increases –> why? 1. the stronger the electrostatic pull (the higher the Zeff) between the nucleus and the valence shell electrons –> 2. the greater the energy released will be when the atom gains an electron
top to bottom = decreases –> why? because the valence shell is farther away from the nucleus as principal quantum number increases
describe the electron affinities of noble gases
EA on the order of 0 because they already possess a stable octet and cannot readily accept an electron
defn: electronegativity
a measure of the attractive force that an atom will exert on an electron in a chemical bond
the greater the EN of an atom, the more it attracts electrons within a bond
how is electronegativity related to ionization energy?
they adjust in parallel
- the lower the ionization energy, the lower the electronegativity
- the higher the ionization energy, the higher the electronegativity
what three elements have negligible electronegativity? why?
the first three noble gases
because they do not form bonds often
range: Pauling electronegativity scale
ranges from 0.7 to 4.0
what is the trend of electronegativity across the periodic table?
L to R: increases
Top to bottom: decreases
SUMMARIZE: periodic trends
Left to Right: atomic radius decreases // ionization energy, electron affinity, and electronegativity increases
Top to bottom: atomic radius increases // ionization energy, electron affinity, and electronegativity decreases
CH 2.1
defn: periodic law
the chemical and physical properties of the elements are dependent, in a periodic way, upon their atomic numbers (number of protons)
defn: period vs. group/family
periods = rows
groups/families = columns
based on atomic number
how do protons and electrons change across a period?
each element in a given period has one more proton and one more electron than the element to its left (in their neutral states)
what is a commonality between elements within in a group (column) of the periodic table?
- same electron configuration in their valence shell
- share similar chemical properties
what does the Roman numeral above each group represent?
the number of valence electrons elements in that group have in their neutral state
defn: A elements vs. B elements
A elements = representative elements = groups IA through VIIIA
- have their valence electrons in the orbitals of either s or p subshells
–> electron configurations determined by the Roman numeral and the letter designation
B elements = nonrepresentative elements = includes transition elements (valence electrons in s and d subshells) and the lanthanide and actinide series (valence electrons in the s and f subshells)
–> may have unexpected electron configurations
CH 2.2
periodic place + what does this group include + char(6): metals
found on the left and middle of periodic table (includes active metals, transition metals, lanthanide and actinide series)
1. lustrous (shiny) (other than mercury)
2. high melting points and densities (other than lithium)
3. malleability
4. ductility (its ability to be pulled or drawn into wires)
5. easily gives up electrons
6. good conductors of heat and electricity (bc the valence electrons of all metals are only loosely held to their atoms, they are free to move)
what are 2 unique characteristics of transition metals (many of the group B elements)?
have two or more oxidation states (charges when forming bonds with other atoms)
some transition metals are relatively nonreactive
char at the atomic level (6): metals
- low effective nuclear charge
- low electronegativity/high electropositivity
- large atomic radius
- small ionic radius
- low ionization energy
- low electron affinity
what subshells are the valence electrons of different metals in?
- active metals = s subshell
- transition metals = s and d subshells
- lanthanide and actinide series = s and f subshells
periodic place + char (5): nonmetals
predominantly on the upper right side of the periodic table
1. brittle when solid
2. show little to no metallic luster
3. poor conductors of heat and electricity
4. inability to easily give up electrons
5. less unified in their chemical and physical properties than the metals
char at the atomic level (5): nonmetals
- high ionization energies
- high electron affinities
- high electronegativities
- small atomic radii
- large ionic radii
periodic place + aka + char (3): metalloids
separate the metals and nonmetals as a stair-step group of elements
aka = semimetals bc they share some characteristics with both metals and nonmetals
- densities, melting points, and boiling points vary widely and can be combos of metallic and nonmetallic characteristics
- reactivities depend on the elements with which they are reacting
- good semiconductors bc of partial conductivity of electricity
char at atomic level: metalloids
electronegativities and ionization energies lie between metals and nonmetals
what are the 8 metalloids that form the staircase?
- boron (B)
- silicon (Si)
- germanium (Ge)
- arsenic (As)
- antimony (Sb)
- tellurium (Te)
- polonium (Po)
- astatine (At)
diagram: periodic table, coded by element type
CH 2.4
group + char (6): alkali metals
Group IA/Group 1
1. have most of the classic physical properties of metals except their densities are lower than those of other metals
2. have only one loosely bound electron in their outermost shells
3. have low Zeff values
4. have the largest atomic radii of all elements in their rows
5. low ionization energy, low electron affinity, low electronegativity
6. easily lose one electron to form univalent cations and react readily with nonmetals (especially the halogens)
group + char (4): alkaline earth metals
group IIA/group 2
1. possess many properties characteristic of metals
2. slightly higher Zeff than alkali metals
3. slightly smaller atomic radii than alkali metals
4. have 2 electrons in their valence shell, which are both easily removed to form divalent cations
what are alkali and alkaline earth metals called together? why?
active metals = they are so reactive that they are not naturally found in their elemental (neutral) state
group + char (6): chalcogens
group VIA/group 16
1. eclectic group of nonmetals and metalloids
2. crucial for normal biological functions
3. each have 6 electrons in their valence shell
4. have small atomic radii
5. large ionic radii
6. at high concentrations, many of these elements can be toxic or damaging, no matter how biologically useful
what are the 3 biologically important elements in the chalcogen group? what are 2 characteristics of the rest of the elements in the group?
- oxygen
- sulfur
- selenium
rest of the group: primarily metallic and generally toxic to living organisms
group + char (8): halogens
what element in this group has the highest electronegativity of all elements?
group VIIA/group 17
1. highly reactive nonmetals with 7 valence electrons
2. desperate to complete their octets by gaining one additional electrons
3. variable physical properties
4. more uniform chemical reactivity
5. high electronegativity
6. high electron affinity
7. very reactive toward alkali and alkaline earth metals
8. so reactive that they are not found in their elemental state, but rather as ions (called halides)
Fluorine has the highest EN of all elements
group + char (5) + aka: noble gases
group VIIIA/group 8
aka inert gases
- minimal chemical reactivity due to their filled valence shells
- high ionization energies
- little or no tendency to gain or lose electrons
- no measurable EN (for at least He, Ne, and Ar)
- extremely low boiling points
- exist as gases at room temperature
group + char (8): transition metals
groups IB to VIIIB/groups 3 - 12
1. considered to be metals
2. low electron affinities, low ionization energies, low electronegativities
3. very hard
4. high melting and boiling points
5. malleable
6. good conductors due to loosely held electrons that progressively fill the d-orbitals in their valence shells
7. many of them can have different possible charged forms/oxidation states bc they are capable of losing different numbers of electrons from the s and d orbitals in their valence shells –> form many different ionic compounds —> different oxidation states often correspond to different colors (solutions with transition metal-containing complexes are often vibrant)
8. tend to associate in solution with molecules of water (hydration complexes) or with nonmetals –> this ability to form complexes adds to the variable solubility of certain transition metal-containing compounds
explain why the formation of complexes allows for the vibrant colors characteristic of transition metals (3)
- the formation of complexes causes the d-orbitals to split into 2 energy sublevels
- this enables many of the complexes to absorb certain frequencies of light (those containing the precise amount of energy required to raise electrons from the lower to the higher energy d orbitals)
- the frequencies not absorbed (subtraction frequencies) give the complexes their characteristic colors
when we perceive an object as a particular color, it is because that color is not absorbed, but rather reflected, by the object. explain.
if an object absorbs a given color of light and reflects all others, our brain mixes these subtraction frequencies and we perceive the complementary color of the frequency that was absorbed
