chemistry quiz 8 (chapter 8 questions 1-40)

define atomic radius. for main-group elements, describe the observed trends in atomic radius as you move: a) across a period in the periodic table b) down a column in the periodic table

1 way to define atomic radii is to consider the distance between nonbonding atoms in molecules or atoms that are touching each other but are not bonded together. atomic radius determined this way is valled the nonbonding atomic radius or the van der Waals radius. the van der waals radius represents the radius of an atom when it is not bonded to another atom another way to define the size of an atom, valled bonding an atomic radius or covalent radius, is defined differently for nonmetals and metals as follows: nonmetals: 1/2 the distance between 2 of the atoms bonded together metal: 1/2 the distance between 2 of the atoms next to each other in a crystal of the metal. A more general term, the atomic radius, refers to a set of average bonding radii determined from measurements on a large number of elements and compounds. the atomic radius represents the radius of an atom when it is bonded to another atom and is always smaller than the van der waals radius. a) as you move to the right across a period, atomic radius decreases b) as move down column in periodic table, atomic radius increases

what is effective nuclear charge? what is shielding?

effective nuclear charge (Zeff)=average or net charge from the nucleus experienced by the e-s in the outermost levels. shielding=the blocking of nuclear charge from the outermost e-s. the shielding is primarily due to the inner (core) e-s, although there is some interaction and shielding from the e- repulsions for the outer e-s with each other.

what is an orbital diagram? provide an example.

orbital diagram =different way to show the e- configuration of an atom. -symbolizes the e- as an arrow in a box that represents the orbital -orbital diagram for hydrogen atom:

what is ionization energy? what is the difference between first ionization energy and second ionization energy?

the ionization energy (IE) of an atom or ion is the energy required to remove an e- from the atom or ion in the gaseous state. the ionization energy is always + because removing an electron takes energy. the energy required to remove the first electron is called the first ionization energy (IE1). the energy required to remove the second e- is called the second ionization energy (IE2). the second IE is always greater than the first IE.

why are the sublevels within a principal level split into different energies for multielectron atoms but not for the hydrogen atom?

the sublevels within a principal level split in multielectron atoms because of penetration of the outer e-s into the region of the core e-s -the sublevels in hydrogen are not split because they are empty in the ground state.

explain the contributions of Johann Dobereiner and John Newlands to the organization of elements according to their properties

-1st attempt to organize elements according to similarities in their properties made by german chemist Johann Dobereiner. grouped elements into triads: 3 elements w similar properties -more complex approach attempted by english chemist john newlands- organized elements into octaves (analogous to musical notes). when arranged this way, the properties of every 8th element were similar.

what are periodic properties?

a periodic property is predictable based on the element's position within the periodic table.

list the number of valence electrons for each family, and explain the relationship between the number of valence electrons and the resulting chemistry of the elements in the family. a) alkali metals b) alkaline earth metals c) halogens d) oxygen family

a) alkali metals (group 1A) have 1 valence e- and are among the most reactive metals because their outer electron configuration (ns^1) is 1 e- beyond a noble gas configuration. they react to lose the ns^1 electron, obtaining a noble gas configuration. this is why the group 1A metals tend to form 1+ cations b) the alkaline earth metals (group2A) have 2 valence e-s, have an outer e- configuration of ns^2, and also tend to be reactive metals. they lose their ns^2 electrons to form 2+ cations. c) the halogens (group 7A) have 7 valence e-s and an outer e- configuration of ns^2np^5. they are among the most reactive nonmetals. they are only 1 e- short of a noble gas configuration and tend to react to gain that 1 e-, forming 1- anions. d) the oxygen family (group 6A) has 6 valence e- and has an outer e- configuration of ns^2np^4. they are two e-s short of a noble gas configuration and tend to react to gain those 2 e-s, forming 2- anions.

describe the relationship between a) the radius of a cation and that of hte atom from which forms b) the radius of an anion and that of the atom from which is forms.

a) in general, cations much smaller than their corresponding parent. this is bc the outermost e-s are shielded from the nuclear charge in the atom and contribute greatly to the size of the atom. when these e-s are removed to form the cation, the same nuclear charge is now acting only on the core e-s b) in general, anions are much larger than their corresponding atoms. this is because the extra electrons are added to the outermost electrons but no additional protons are added to increas the nuclear charge. the extra electrons increase the repulsions among the outermost electrons, resulting in an anion that is larger than the atom.

for transition elements, describe and explain the observed trends in atomic radius as you move: a) across a period in the periodic table b) down a column in the periodic table.

a) radii of transition elements stay roughly constant across each row instead of decreasing in size as in the main-group elements -the difference is that across a row of transition elements, the number of e-s in the outermost principal energy level is nearly constant. -as another proton is added to the nucleus with each successive element, another e- is added as well, but the e- goes into an h(highest)-1 orbital -the number of outermost e-s stays constant, and they experience a roughly constant effective nuclear charge, keeping the radius approximately constant b) as you go down the first 2 rows of a column within the transition metals, the elements follow the same general trend in atomic radii and the main-group elements; that is, the radii get larger because you are adding outermost e-s into higher n levels.

write a general equation for the reaction of an alkali metal with each substance a) a halogen b) water

a) the reactions of the alkali metals with halogens 2 M(s) + X2 --> 2MX(s) b) alkali metals react with water to form the dissolved alkali metal ion, the hydroxide ion, and hydrogen gas: 2 M(s) + 2H2O(l) --> 2M+(aq)+2OH-(aq)+H2(g)

write a general equation for the reaction of a halogen with each substance a) a metal b) hydrogen c) another halogen

all of the halogens are powerful oxidizing agents a) halogens react with metals to form metal halides: 2M(s)+nX2 --> 2MXn(s) b)the halogens react with hydrogen to form hydrogen halides: H2(g)+X2 --> 2HX(g) c) the halogens react with each other to form interhalogen compounds. for ex: Br2(l)+F2(g) -->2BrF(g)

what is an electron configuration? give an example

an electron configuration shows the particular orbitals that are occupied by electrons in an atom. Some examples are H=1s^1, He=1s^2, Li=1s^22s^1

describe how to write an electron configuration for a transition metal cation. is the order of electron removal upon ionization simply the reverse of electron addition upon filling? why or why not?

an important exception to simply subtracting the number of electrons occurs for transition metal cations. when writing the e- configuration of a transition metal cation, remove the e-s in the highest n-value orbitals first, even if this does not correspond to the reverse order of filling. normally, even though the d orbital electrons add after the s orbital electrons, the s orbital electrons are lost first. this is because (1) the ns and (n-1)d orbitals are extremely close in energy and, depending on the exact configuration, can vary in relative ordering and (2) as the (n-1)d orbitals begin to fill in the first transition series, the increasing nuclear charge stabilizes the (n-1)d orbitals relative to the ns orbitals. this happens because the (n-1)d orbitals are not outermost orbitals and therefore are not effectively shielded from the increasing nuclear charge by the ns orbitals.

use the concepts of effective nuclear charge, shielding, and n value of the valence orbital to explain the trend in atomic radius as you move across a period in the periodic table.

as you move to the right across a row in the periodic table, the n level stays the same. but nuclear charge increases and the amount of shielding stays about the same because the number of inner e-s stays the same. -so the effective nuclear charge experienced by the e-s in the outermost principal energy level increases, resulting in a stronger attraction between the outermost e-s and the nucleus and, therefore, smaller atomic radii.

what is coulomb's law? explain how the PE of 2 charged particles depends on the distance between the charged particles and on the magnitude and sign of their charges.

coulomb's law states that the potential energy (E) of two charged particles depends on their charges (q1 and q2) and on their separation,r. E=(1/4piEo)(q1q2/r). -PE is + for charges of the same sign and - for charges of opposite sign -magnitude of the PE depends inversely on the separation between the charged particles

what are degenerate orbitals? according to hund's rule, how are degenerate orbitals occupied?

degenerate orbitals=same energy. in multielectron atom, orbitals in a sublevel are degenerate. hund's rule=when filling degenerate orbitals, e-s fill them singly first w parallel spins. this is a result of an atom's tendency to find the lowest energy state possible.

what are the exceptions to the periodic trends in first ionization energy? why do they occur?

exceptions occur w elements Be, Mg, and Ca in group 2A having a higher first ionization energy than elements B, Al, and Ga in group 3A. this exception is caused by the change in going from the s block to the p block. the result is that the e-s in the s orbital shield the e- in the p orbital from nuclear charge, making it easier to remove another exception occurs, with N, P and As in group 5A having a higher first ionization energy that O, S, and Se in group 6A. this exception is caused by the repulsion between e-s when they must occupy the same orbital. group 5A has 3p e-s whereas group 6A has 4p electrons. in teh group 5A elements, the p orbitals are half-filled, which makes the configuration particularly stable. the 4th group 6A e- must pair with another e-, making it easier to remove.

list all orbitals from 1s through 5s according to increasing energy for multielectron atoms.

in order of increasing energy the orbitals are 1s<2s<2p<3s<3p<4s<3d<4p<5s 4s orbital fills before the 3d and the 5s fills before the 4d. they are lower in energy because of the greater penetration of the 4s and 5s orbitals.

which of the transition elements in the first transition series have anomalous electron configurations?

in the first transition series of the d block, Cr and Cu have anomalous e- configurations. Cr expected to be [Ar]4s^23d^4, but is found to be [Ar]4s^13d^5, and Cu is expected to be [Ar]4s^23d^9, but is found to be [Ar]4s^13d^10.

what is the general trend in first ionization energy as you move down a column in the periodic table? as you move across a row?

ionization energy generally decreases as you move down a column in the periodic table because electrons in the outermost principal level become farther from the positively charged nucleus and are therefore held less tightly. ionization energy generally increases as you move to the right across a period in the periodic table because electrons in the outermost principal energy level generally experience a greater effective nuclear charge; therefore, the electrons are closer to the nucleus.

what is metallic character? what are the observed periodic trends in metallic character?

metals are good conductors of heat and electricity, they can be pounded into flat sheets (malleability), they can be drawn into wires (ductility), they are often shiny, and they tend to lose e-s in chemical reactions. as you move to the right across a period in the periodic table, metallic character decreases. as you move down a column in the periodic table, metallic character increases

explain the contributions of Meyer and Moseley to the periodic table.

meter proposed an organization of the known elements based on some periodic properties -moseley-listed elements according to the atomic number rather than the atomic mass. resolved the problem's in mendeleev's table where an increase in atomic mass did not correlate with similar properties.

who is credited with arranging the periodic table? how are the elements arranged in the modern periodic table?

modern periodic table credited primarily to Russian chemist Dmitri Mendeleev. his table based on periodic law-when elements arranged in order of increasing mass, their properties recur periodically. -arranged elements in a table in which mass increased from left to right and elements w similar properties fell in the same columns.

why is electron spin important when writing electron configurations? explain in terms of the pauli exclusion principle

pauli exclusion principle: no 2 e-s in an atom can have the same 4 quantum numbers. -because 2 e-s occupying the same orbital have 3 identical quantum numbers (n, l, ml), they must have different spin quantum numbers -pauli exclusion principle implies that each orbital can have a maximum of only 2 e-s w opposing spins.

what is penetration? how does the penetration of an orbital into the region occupied by core electrons affect the energy of an electron in that orbital?

penetration-occurs when an e- penetrates the e- cloud of the 1s orbital and expEriences the charge of the nucleus more fulLY bc it is less shielded by the intervening e-s. -as the outer e- undergoes penetration into the region occupied by the inner e-s, experiences a greater nuclear charge and therefore, according to coulomb's law, a lower energy

the periodic table is a result of the periodic law. what observations led to the periodic law? what theory explains the underlying reasons for the periodic law?

periodic law based on the observations that the properties of elements recur and certain elements have similar properties -theory that explains the existence of the periodic law=quantum-mechanical theory

which periodic property is particularly important to nerve signal transmission? why?

relative sizeof sodium and potassium ions imp to nerve signal transmission -pumps and channels within cell membranes so sensitive that can distinguish between sizes of these two ions and selectively allow only 1 or the other to pass -movement of ions=basis for the transmission of nerve signals in the brain and throughout the body.

why do the rows in the periodic table get progressively longer as you move down the table? for example, the first row contains 2 elements, the second and third rows each contain 8 elements, and the fourth and fifth rows each contain 18 elements. explain.

rows in periodic table grow progressively longer because you are adding sublevels as the n level increases

copy a blank periodic table onto a sheet of paper and label each of the blocks within the table: s block, p block, d block, and f block

s block elements=first 2 columns + He d block elements=transition metals p block elements=last 6 columns f block elements=lanthanides and actinides, bottom 2 rows below periodic table.

what is shielding? in an atom, which e-s tend to do the most shielding (core electrons or valence electrons?)

shielding/screening occurs when 1 e- is blocked from the full effects of the nuclear charge so that the e- experiences onlY a part of the nuclear charge. -it is the inner/core e-s that shield the outer e-s from the full nuclear charge

describe the relationship between the properties of an element and the number of valence electrons that it contains.

the chemical properties of elements largely determined by the number of valence e-s they contain -properties periodic bc the number of valence e-s is periodic. -because elements within a column in the periodic table have the same number of valence electrons, they also have similar chemical properties.

what is electron affinity? what are the observed periodic trends in electron affinity?

the e- affinity (EA) of an atom or ion is the energy change associated with the gaining of an e- by the atom in the gaseous state. the e- affinity is usually-though not always- negative because an atom or ion usually releases energy when it gains an e-. the trends in EA are not as regular as trends in other properties. for main-group elements, EA generally becomes more negative as you move to the right across a row in the periodic table. there is no corresponding trend in EA going down a column, with the exception of group IA which becomes more positive as you go down the column

how is the electron configuration of an anion different from that of the corresponding neutral atom? how is the electron configuration of a cation different?

the electron configuration of a main-group monatomic ion can be deduced from the e- configuration of the neutral atom and the charge of the ion. for anions, add the number e-s required by the magnitude of the charge of the anion. the e- configuration of cations is obtained by subtracting the number of e-s required by the magnitude of the charge.

describe the relationship between a main group element's lettered group number (the number of the element's column) and its valence electrons.

the lettered group number of a main-group element is = to the number of valence e-s for that element.

explain why the s block in the periodic table only has 2 columns while the p block has 6

the number of columns in a block corresponds to the maximum number of e-s that can occupy the particular sublevel of that block. -s block has 2 columns corresponding to one s orbital holding a maximum of 2 e-s. -p block has 6 columns corresponding to the 3 p orbitals with 2 e-s each

describe the relationship between an element's row number in the periodic table and the highest principal quantum number in the element's electron configuration. how does this relationship differ for main-group elements, transition elements, and inner transition elements?

the row number of a main-group element is = to the highest principal quantum number of that element. however, the principle quantum number of the d orbital being filled across each row in the transition series is = to the row number -1. for the inner transition elements, the principal quantum number of the f orbital being filled across each row is the row number minus 2

examination of the first few successive ionization energies for a given element usually reveals a large jump between two ionization energies. For example, the successive ionization energies of magnesium show a large jump between IE2 and IE3. the successive ionizaton energies of aluminum show a large jump between IE3 and IE4. explain why these jumps occur an how you might predict them.

the second ionization energy of Mg involves removing the second outermost e- leading to an ion with a noble gas configuration for hte core e-s. the third ionization energy requires removing a core e- from an ion with a noble gas configuration. this requires a tremendous amount of energy, making IE3 very high. for Al, IE3 involves removing the third outermost e- for Al, leaving the ion with a noble gas configuration of the core e-s. IE4 then requires removing a core e- from an ion with a noble gas configuration. this requires a tremendous amount of energy and makes IE4 very high you can predict whether the IE energy is going to be very high by looking for the ionization that requires removing a core electron.

describe how to write the electron configuration for an element based on its position in the periodic table.

to use the periodic table to write the e- configuration, find the noble gas that precedes the element. -element has the inner e- configuration of that noble gas -place symbol for the noble gas in [ ]. obtain the outer e- configuration by tracing the element across the period and assigning e=s in the appropriate orbitals.

what are valence electrons? why are they important?

valence e-s imp in chemical bonding. for main-group elements, valence e-s in outermost principal energy level. -for transition elements, we also count the outermost d electrons among the valence, even though they are not in the outermost principal energy level -chemical properties of an element depend on its valence e-s, which are imp in bonding because they are held most loosely. this is why the elements in a column of the periodic table have similar chemical properties; they have the same number valence e-s.