Chem exam 3
Formula fr 2nd ionization energy
I,<I2,<I3<...
Octet rule (for atoms of main group elements only)
Atoms lose or gain electrons to form ions (cations and anions) whose configurations resemble those for the closest noble gas: ns^2np^6 or ns^2 (n-1)d^10 np^6
What would be isoelectronic with sulphide?
Chloride, potassium, calcium.. etc
True or false: The hydrogen atom is most stable when it has a full octet of electrons. false true
False
The ________ ion is represented by the electron configuration [Ar]3d2.
Fe^6+
Which atom do you expect to gladly accept an e^-?
Group VIIA (7A) w/ 7 valence electrons, so that octet can be achieved
Electron affinity, EA, = ?
How much energy is released (EA<0) or gained (EA >0) when an atom gains an electron
Does metallic character increase, decrease, or remain unchanged as one goes down a column of the periodic table?
Increase
What is I1?
Ionization energy
As zeff decreases what happens to Ra?
It increases
The following ions contain the same number of electrons. Rank them in order of decreasing ionic radii.
N3-, O2-, F-, Na+,Mg2+, Al3+
Does the octet rule apply to transition metals?
No
Do noble gases intermingle with other elements?
No they're perfect and wanna stay that way
Main group elements : 11Na: 1s^2 2s^2 2p^6 3s^1 ---->
So we need to lose one electron. It will look like: Na^+: 1s^2 2s^2 2p^6 **Na is in group 1A and adopts the charge equal to its group #. *1s^2 2s^2 2p^6 is the preceding noble gas configuration
In which orbital does an electron in a nitrogen atom experience the greatest shielding?
The 2p sub-level would experience the greatest shielding - there are 4 intervening electrons between the nucleus and the 3 2p electrons in nitrogen.
What does ionization mean?
The First Ionization of an atom means that you remove one electron; the second ionization means that you are removing two electrons from the atom; etc.
What does isoelectronic mean?
The ions have the same total number of electrons (and same electron configuration)
Formula for atom in ground state
(g) + e^- -----> -1 anion (g) ; DeltaE=EA.
What is the most common charge of transition metals?
+2
What is the most likely charge of zinc in an ionic compound? {Ar] 4s^2 3d^10
+2
What situations do we want to avoid in Lewis structures?
- Adjacent atoms that have the same sign f.ch's - The more electronegative atom in a bond would have the more positive f.ch than the less electronegative atom
Formal charge -Provides: -Is calculated for:
- Provides a way of determining which lewis structure is the best - It is calculated for each individual atom within a molecule or ion
In which orbital does an electron in a bromine atom experience the greatest effective nuclear charge? 4p 4s 3d 1s 4d
1s
Fe2+ ions are represented by the electron configuration ________. [Ar]3d4 [Ar]3d6 [Ar]4s23d4 [Kr]3d6 [Ar]3d8
2nd
Which electrons in a calcium atom belong to the valence? 1s^2 2s^2 2p^6 3s^2 3p^6 4s^2
4s^2 | | -2e^- | | Ca:^2+:[Ne] 3s^2 3p^6
A tin atom has 50 electrons. Electrons in the ________ subshell experience the lowest effective nuclear charge. A tin atom has 50 electrons. Electrons in the ________ subshell experience the lowest effective nuclear charge. 5s 5p 1s 3p 3d
5p
Which atom in any given period (row) do you expect to hang on to its electron and have the highest I1?
The right one, the smallest one at the end of the period (noble gasses)
Why do atoms get bigger as you go down the group and smaller as you go across?
They get smaller because the less electrons b/w the valence electrons and nucleus, the greater the nuclear charge. This means that the electron will be sucked closer to the nucleus making the atom smaller
Which lewis structure is the most stable?
Those where an individual atoms' formal charge is as close to 0 as possible.
A Lewis structure is:
a graphic representation of the valence electron distribution within COVALENTLY bonded molecules or ions. There may be more than one correct lewis structure per species
Nonmetal oxides, when dissolved in water, form ______
acids
_______electrons screen ________ electrons
core valence
As the electron-nucleus attraction increases, ionic size ________
decreases
The effective nuclear charge of an atom is primarily affected by ________. inner electrons orbital radial probability nuclear charge outer electrons electron distribution
inner electrons
What happens during screening (shielding)?
inner electrons screen outer electrons
Anions are ________than their parent atoms
larger
When forming anions of nonmetals, simply add enough electrons to fill the ______
octet
Cations are ______ than their parent atoms
smaller
The further up and to the right of the periodic table, the _________ the atomic radius (generally)
smaller
When an atom loses electrons the resulting cation is ________, even though the nucleus and protons remain unchanged
smaller
Bond dissociation Enthalpy (D)=
the amount of energy in kJ/mol needed to break 1 mole of bonds of a particular type
True or false: The hydrogen atom is most stable when it has a full octet of electrons.
False An octet requires eight electrons in the valence shell. Hydrogen's valence shell is the n=1 shell, which can carry only two electrons at most. This is the 1s subshell, with one orbital.
What is the second ionization energy (I2)?
It is how much energy is necessary to remove a 2nd electron
Are the periodic trends observed in Parts A and B the same as or different from those for first ionization energy? The trends in Parts A and B are equal to the trends in ionization energy. The trend in Part A is equal to and the trend in Part B is opposite the trends in ionization energy. The trends in Parts A and B are the opposite of the trends in ionization energy. The trend in Part A is opposite and the trend in Part B is equal to the trends in ionization energy.
The trends are opposite
The more electrons, the _______ an atom gets
bigger
Screenin (shielding) of electrons in regard to quantum numbers
s<p<d<f The higher the quantum number, the more screening will take place.
Penetration:
s>p>d>f
For electrons in s msn electron atom: within a shell (same n) l goes up: what goes up and what goes down?
screening goes up penetration goes down Zeff goes down Attraction b/w electrons and nucleus goes down Eelectron goes up
Covalent bonding The octet is_______
still achieved, but no ions are formed
Bonds between small atoms tend to be_______ than between large atoms
stronger
LiF has a _________lattice than KI
stronger Li and K both have +1 charge. F and I both have -1 charge, so ionic charge won't contribute to determining which one is stronger. The only thing left is to look at size. Li is smaller than K and F is smaller than I
First ionization energy; definition and equation
the first ionization energy is the energy required to remove one mole of the most loosely held electrons from one mole of gaseous atoms to produce 1 mole of gaseous ions each with a charge of 1+ . The equation below represents removing an electron from each atom in a mole of Ne atoms: Ne (g) ⟶ Ne⁺ (g) + e⁻ first ionization energy ∆H⁰ = 2080.6 kJ/mol
F.ch = ?
# of valence e^-1's - (# of lone pair e^-1s + 1/2 # of bonding e^-1s
Anions are _____ the their parent atoms
larger
Which atom in group 1A do you expect to lose its one electron the easiest? (have the lowest I1)
the largest, where e^- is the furthest from the nucleus, Cs (or Fr)
**Skipped a question. Identify bromines valence electrons
1s^2 2s^2 2p^6 3s^2 3p^6 4s^2 3d^10 4p^5 4s^2 4p^5
The halogens, alkali metals, and alkaline earth metals have ________ valence electrons, respectively.
7, 1 and 2
Basically: cations are smaller than their parent atoms. Why?
Because there are less electrons to attract so the ones there can be pulled closer to the nucleus
How do you determine which atoms possess a higher ionization energy?
Draw their electron configuration!
Which of the following correctly lists the five atoms in order of increasing size (smallest to largest)? F < Ge < Br < K < Rb F < Br < Ge < Rb < K F < K < Ge < Br < Rb F < Br < Ge < K < Rb F < K < Br < Ge < Rb
E
Electronegativity will follow the same periodic trends as ____
I1 and EA
Rank the following items in order of decreasing radius: Mg, Mg2+, and Mg2−.
Mg2-, Mg, Mg2+
An outer electron (in an orbital further from the nucleus) experiences an effective nuclear charge of _________
Zeff<<Z
Many metal; oxides as well as peroxides and superoxides, when dissolved in water, form ________
bases
Does metallic character increase, decrease, or remain unchanged as one goes from left to right across a row of the periodic table? increase decrease remain unchanged
decrease
The more negative the isoelectric is the _______ it is
larger
The majority of bases are _____
metal hydroxides: (Li)OH, (Ca)(OH)2, (Al)(OH)3
MgO has _________lattice than NaCl
stronger Mg form +2 and NA is +1. Mg has ore positive, more protons, so it can pull electrons closer to its nucleus , making Mg smaller. O is smaller than Cl as well
Electrons within a _________ do not screen each other very well
subshell
Covalent bonding occurs;
when two atoms share electrons due to a partial transfer (shift) of electrons from one atom to another
For the series of elements X, Y, and Z all in the same period (row), arrange the elements in order of decreasing first ionization energy. Element Radius (pm) X 109 Y 178 Z 272
x y z
Zeff equation
zeff = Z-S, where Z= the number of protons and S=the screening constant. *Basically, the atomic number minus the electrons that aren't the valence
ionization energy formula:
atom in ground state (g) + energy ---> +1 cation (g) +e^- *I1> 0 because energy goes into it
Ionic bonding (2 things)
- Electrostatic attraction between oppositely charged ions - Occurs when there is a COMPLETE TRANSFER of electrons from one atom to another
What does bond strength depend on?
- The number of electrons in a bond (related to bond order) - The sizes of atoms connected via the bond - The electronegativities of atoms in a bond
Using the table of bond dissociation energies, the ΔH for the following reaction is ________ kJ. 2HCl (g) + F2 (g) → 2HF (g) + Cl2 (g) BondH−ClF−FH−FCl−ClD(kJ/mol)431155567242 Using the table of bond dissociation energies, the H for the following reaction is ________ kJ. 208 223 -359 -223 359
-359
Using the table of bond dissociation energies, the ΔH for the following gas-phase reaction is ________ kJ. There is a scheme of an irreversible chemical reaction. Here, the reactants are C2H4 and HBr, and the product is C2H5Br. The C2H4 molecule is two C connected with a double bond and each of them has 2 H atoms attached. HBr is H connected to Br by a single bond. And C2H5Br is an HCCBr molecule with 2 H atoms attached to each carbon. BondC−CC=CC−HH−BrC−BrD(kJ/mol)348614413366276 Using the table of bond dissociation energies, the H for the following gas-phase reaction is ________ kJ. -57 291 2017 -291 -356
-57
Consider the following oxides: SO2, Y2O3, MgO, Cl2O, and N2O5. How many are expected to form acidic solutions in water?
3 Metal oxides react with water to produce basic solutions. Nonmetal oxides react with water to produce acidic solutions. SO2(g) + H2O(l) <==> H+ + HSO3^- ..... acidic Y2O3(s) + 3H2O(l) <==> 2Y(OH)3(s) .... basic MgO(s) + H2O(l) <==> Mg(OH)2(s) ....... basic Cl2O(g) + H2O(l) <==> 2HOCl(aq) ........ acidic N2O5(s) + H2O(l) <==> 2HNO3(aq) ...... acidic
In the ground-state electron configuration of Fe3+, how many unpaired electrons are present?
5 Iron as an atom has 4 electrons that are unpaired: 3D (II) (I ) (I) (I )(I ), 4S(II) But it loses the 2 electrons in the 4th S & 1 electron in the 3D to become Fe+3: 3D (I ) (I ) (I) (I )(I ), 4S( ) therefor all of the 3rd shell D- orbitals are unpaired
Main group elements 19In: {Kr} 5s^2 4d^10 5p^1 (this is the electron we will lose) ----->
5p^1 (- 1e^-) -----> In^+: [Kr] 5s^2 4d^10 (-2e^-) ---> In:^3+: [Kr] 4d^10 *Note that In is in group 3
How many dots does a lewis symbol of sulfur have?
6
What is the total number of valence electrons in SCN^-
6+4+%+1= 16
Rank the following ions in order of decreasing radius: Be2+,Mg2+,Ca2+,Sr2+, and Ba2+.
Ba, Sr, Ca, Mg, Be **Radius increases as you go down and to the left
Why does atomic radii increase going down a group?
Because successively larger valence-shell orbitals are occupied by electrons. For example, iodine has electrons in the fifth shell, which contains much larger orbitals than the fourth, third, second, or first shells.
Why do atoms get smaller as you go -----> across the periodic table?
Because we're adding electrons and that means more protons. The electrons aren't screened as welll, pulling atoms closer to the nucleus and making the atom's size smaller. As you move from left to right across a period, the number of protons in the nucleus increases. The electrons are thus attracted to the nucleus more strongly, and the atomic radius is smaller (this attraction is much stronger than the relatively weak repulsion between electrons). As you move down a column, there are more protons, but there are also more complete energy levels below the valence electrons. These lower energy levels shield the valence electrons from the attractive effects of the atom's nucleus, so the atomic radius gets larger.
There is a big sphere, a medium one, and a small one. Which spheres would the charges F, Br, and Br^- represent?
Big=Br Medium- Br- Small=F
Which of these elements is most likely to form ions with a 2+ charge? P Li Ca O Cl
Ca Calcium is found in the second group of the periodic table and it is most likely to lose two electrons to form a 2+ charged cation, Ca2+.
Bond dissociation enthalpies have positive values. What kind of reactions are these?
Endothermic. To break a bond you'll have to work for it
Write equation for the electron affinity of fluorine.
F(g) + e^- ---> F^- (g)
Which group of atoms do you expect to most easily lose their one electron?
First main group, group 1A. These are the alkali metals: Li, Na, K... The largest atoms within each period
Arrange the elements in decreasing order of first ionization energy Cs,Ge,Ln,Se
Here: decreasing 1st ionization energy is APPROXimately inversely related to atomic radius: ioniz275 < ioniz185 < ioniz112 Listed in decreasing 1st ionization (from source) Se 9.752 (radius 1.4) Ge 7.899 (radius 1.37) In 5.786 (radius 1.66) Cs 3.894 (radius 2.67)
A sample of soil from a newly discovered cave is analyzed by a team of explorers. They find an element that is a good conductor of electricity. It also forms a chloride in the form XCl2 and an oxide in the form XO. The element is a liquid at room temperature. What is the identity of this element?
Hg
The shielding of electrons gives rise to an effective nuclear charge, Zeff, which explains why boron is larger than oxygen. Estimate the approximate Zeff felt by a valence electron of boron and oxygen, respectively?
Hints: How many protons and how many inner (nonvalence) electrons does a neutral boron atom have? 5,2 How many protons and how many inner (nonvalence) electrons does a neutral oxygen atom have? 8,2 answer is 3 and 6. The valence electrons in an oxygen atom are attracted to the nucleus by a positive charge nearly double that of boron. Therefore, the electrons in oxygen are held closer to the nucleus, giving it a smaller radius. **Larger the zeff, the smaller the atom
Suppose that a metal oxide of formula M2O3 were soluble in water. What would be the major product or products of dissolving the substance in water? M(OH)3(aq) M3+(aq) + H2O2(aq) MH3(aq) + O2(g) M(OH)2(aq) M(s) + H2(g) + O2(g)
M(OH) The reaction between a soluble metal oxide and water produces the metal hydroxide which in this case is M(OH)3(aq). Consider a metal oxide with the formula M2O3. Such a compound is likely to be quite insoluble in wter When combined with water some of it may hydrolyze (react with water) to make the insoluble metal hydroxide. Since there is no oxidation or reduction involved, the metal will retain its +3 oxidation state and the formula will be M(OH)3. M2O3(s) + 3H2O(l) --> 2M(OH)3(s)
Octet rule:
Main group elements tends to lose or gain (or share) electrons until they possess 8 valence electrons
How do you rank ions (atoms with different charges the usual) from biggest to smallest?
Make a list of their protons and electrons. If isoelectronic, the ones with the highest positive force will be the largest
Which of these elements is unlikely to form covalent bonds? S, H, Mg, Ar, Si S H Mg Ar Si
Mg, Ar Mg is a group 2A (2) metal with two valence electrons, likely to form an ion with a 2+ charge and participate in ionic bonding to achieve a full octet of electrons. Ar is in group 8A (noble gases), which means it has a full octet of electrons and so is not likely to form any type of bond.
Does the number of protons in the nucleus of an atom change when an atom becomes an ion?
No, only the electrons change
Do hydrated ions follow the same ionic size trend that naked ions do?
No- it has water molecules attached to it
Do transition metals always follow the octet rule?
No- transition metals can have different fillings of d sub shell, think Roman numerals. They can have varying charges
How many electrons can be put into the 3px orbital?
Only 2 ___________ _________ _________ x y z all makes up the 3p
Consider the following properties of an element: (i) It is solid at room temperature. (ii) It easily forms an oxide when exposed to air. (iii) When it reacts with water, hydrogen gas evolves. (iv) It must be stored submerged in oil. Which element fits the above description the best? sodium copper mercury sulfur magnesium
Sodium
Which of the following would have to lose two electrons in order to achieve a noble gas electron configuration? O Sr Na Se Br O Sr Na Se Br Sr Br Na O, Se Sr, O, Se
Sr
Slaters rule
The general principle behind Slater's Rule is that the actual charge felt by an electron is equal to what you'd expect the charge to be from a certain number of protons, but minus a certain amount of charge from other electrons. Slater's rules allow you to estimate the effective nuclear charge Zeff from the real number of protons in the nucleus and the effective shielding of electrons in each orbital "shell" (e.g., to compare the effective nuclear charge and shielding 3d and 4s in transition metals). Slater's rules are fairly simple and produce fairly accurate predictions of things like the electron configurations and ionization energies. Step 1: Write the electron configuration of the atom in the following form: (1s) (2s, 2p) (3s, 3p) (3d) (4s, 4p) (4d) (4f) (5s, 5p) . . . Step 2: Identify the electron of interest, and ignore all electrons in higher groups (to the right in the list from Step 1). These do not shield electrons in lower groups Step 3: Slater's Rules is now broken into two cases: the shielding experienced by an s- or p- electron, electrons within same group shield 0.35, except the 1s which shield 0.30 electrons within the n-1 group shield 0.85 electrons within the n-2 or lower groups shield 1.00 the shielding experienced by nd or nf valence electrons electrons within same group shield 0.35 electrons within the lower groups shield 1.00
Which of the following statements about effective nuclear charge for the outermost valence electron of an atom is incorrect? - Effective nuclear charge increases going left to right across a row of the periodic table. - The change in effective nuclear charge going down a column of the periodic table is generally less than that going across a row of the periodic table. - Valence electrons screen the nuclear charge more effectively than do core electrons. - The effective nuclear charge shows a sudden decrease when we go from the end of one row to the beginning of the next row of the periodic table. - The effective nuclear charge can be thought of as the true nuclear charge minus a screening constant due to the other electrons in the atom.
Valence electrons scree..
Electronegativity symbols
X or EN
Is it gonna be harder to remove a 2nd electron than a first? Why?
Yes because it will have lower energy. Also, the same nuclear charge is attracting fewer electrons
If the core electrons were totally effective at screening the valence electrons and the valence electrons provided no screening for each other, what would be the effective nuclear charge acting on the 3s and 3p valence electrons in P?.
Z eff = 5 Atomic # for P is 15. Its electron configuration is 1s^2 2s^2 2p^6 3s^2 3p^3. It can be written as [Ne] 3s^2 3p^3. This means that there are 10 core electrons Calculation of Zeff for 3s electrons =15-10X1 =15-10 =5
Electronegativity is the:
ability of an atom to attract shared electrons
In the group 3 to group 12 elements, which subshell is filled up going across the rows?
d sub shell In the group 3 to group 12 transition metals, the outermost s electron shell contains one or two electrons. However, in these metals, it is the d subshells that fill up going across the row. In period 4 of the table, the 3d subshell fills, and in periods 5 and 6, the 4d and 5d subshells fill, respectively. It is important to keep in mind that this filling is not always regular. For example, in period 4, element 23, vanadium, has an electron configuration of [Ar]3d34s2, but element 24, chromium, has an electron configuration of [Ar]3d54s.
The probability of finding electrons belonging to a given orbital close to the nucleus, at least some of the time, _______ with ________. In other words, the higher the _________ the further away the electrons tend to be
decreases l l
Consider the addition of an electron to the following atoms from the third period. Rank the atoms in order from the most negative to the least negative electron affinity values based on their electron configurations. Atom or ion Electron configuration Cl 1s22s22p63s23p5 Si 1s22s22p63s23p2 Ar 1s22s22p63s23p6
electron affinitiy --> the nearest to a noble gas, the stronger note that for a noble gas, electron affinity = 0 by definition Br has a higher electron affinity (wants to get an extra electron so STRONG, so it can form a noble gas config) Ge has medium electron affinity, since it must get 4 electrons or get rid of 4 electrons, cl-si-ar
Define lattice energies;
energy required to completely separate a mole of a solid ion into its gaseous form
Bond dissociation energies can be used to approximate ________ taking place during chemical reactions in which bonds are being broken (in reactants) and formed (in products)
enthalpy changes (delta H)
For main group elements, the # of valence electrons =
group #
Which of the following shows the correct order of increasing lattice energy for the given compounds? Na2O < K2O < Li2O Na2O < Li2O < K2O Li2O < Na2O < K2O K2O < Na2O < Li2O
last
We say that within a shell (the same valence of n for all electrons), s electrons are screened ______than p electrons, p electrons are screened ______ than d electrons...
less
Screening of the nuclear charge by core electrons in atoms is ________. - responsible for a general decrease in atomic radius going down a group - more efficient than that by valence electrons - essentially identical to that by valence electrons - less efficient than that by valence electrons - both essentially identical to that by valence electrons and responsible for a general decrease in atomic radius going down a group
more efficient
How does atomic radii increase? (periods)
moving left to right across the periodic table, atomic radii decreases (Li has larger radius than Be)
Trends in EA are not as clear as others but in general, nonmetals have more _________ EA than metals
negative (more exothermic)
The majority of acids contain a _______
nonmetal : H2(S)O4, H2(C)O3, H(N)O2
Covalent bonding is common between:
nonmetals
Group # General valence Electron configuration IA
ns^1
Group # General valence Electron configuration IIA
ns^2
Group # General valence Electron configuration VA
ns^2np3
Group # General valence Electron configuration IIIA
ns^2np^1
Group # General valence Electron configuration IVA
ns^2np^2
Group # General valence Electron configuration VIA
ns^2np^4
Group # General valence Electron configuration VIIA
ns^2np^5
Group # General valence Electron configuration VIIIA
ns^2np^^
When forming cations of main group metals, first remove electrons from the ______subshell the from the ______ subshell
p s
Charges of ions are determined by electron configurations of their _________, namely the # of __________
parent atoms valece electrons
Cations are ________ charged Anions?
positively negatively
In isoelectric ions, the number of _____ = the number of _______-
protons in the nucleus electrons
The smaller the atom, the harder it will be to ______the electron, the higher the ________
remove ionization energy
We say that ______ electrons penetrate towards the nucleus better than _______ electrons
s p
Recall that L = 0, 1, 2, 3
s p d f
When forming cations of transition metals, first remove electrons from the _________, then (if needed) from the _________
s subshell d subshell