Electron Affinity the Energy Associated with Adding an Electron to an Atom

UMass Boston, Chem 115

CHEM 115 Periodic Trends (Ch. 7), and The Lewis Structure Model (Ch. 8)

Lectures 21 and 22 Prof. Sevian

April 22 is Earth Day. In honor of it, I offer some suggestions for ways to reduce your carbon footprint.

Reduce paper usage. Print on the back of scrap paper when you need to print. Agenda Chapter 7 z Periodic properties z Ionization energy z Atomic radius z Others z Interpreting measured properties of elements in light of their electronic configurations z (+) Core = nucleus + (all but the outer shell of electrons) z (-)V) Val ence = the ou termost sh ell o f e lec trons z Effective nuclear charge = (Total electrons) – (Core electrons) = Zeff z Coulombic force of attraction between core (+) and valence (-) z Building a logical explanation

Convert monthly bills over to automated electronic bills What We’ve Observed So Far instead of bills mailed to you. z Noble gases (Group VIIIA) z High ionization energies compared to other Groups z Within same group, ionization energy is smaller as atomic number increases z All elements in ggproup have com plete shell electron confi gurations (outermost s- and p-subshells are completely filled) z Alkali metals (Group IA) z Low ionization energies compared to other Groups z Within same group, ionization energy is smaller as atomic number increases z All elements in group have one electron in outermost s-subshell z +1 charged ions of Alkali metals (Group IA) z The alkali metals have the same second ionization energy behavior as noble gases z Electron configurations same as noble gases z Trends down a group all follow the same pattern z The trend in the property is that it increases or decreases as you go down the group z The reason is that the most loosely bound electron is less tightly held to the atom

These are all trends that occur within the same group. What about trends across a period? To be able to explain further, we need Coulomb’s law.

Zeff = Z – S Core vs. Valence S = screening constant S is approximately equal to the One way to view an atom or ion number of core electrons (Note: not drawn to scale!)

Nucleus Book calls (protons and this Zeff neutrons) CORE + All complete charged inner shells of electrons

Outer electrons VALENCE beyond all SHELL complete inner – shells charged Image modified from http://www.badastronomy.com/bad/movies/thecore_review.html

Zeff = Z – S = 11 – 10 = 1 Example: Sodium Atom Z = 11 (the atomic number) S = approximately equal to the A neutral sodium atom has 11 protons number of core electrons = 10 and 11 electrons Electronic configuration is 1s2 2s2 2p6 3s1 (Note: not dra wn to scale!)

11 protons and Book calls some neutrons this Z eff (charge: +11) CORE All complete Net +1 inner shells of charge electrons ((gcharge: -10) 1s2 2s2 2p6 3s1 Outer electrons VALENCE beyond all SHELL complete inner Net -1 shells (charge: -1) charge Image modified from http://www.badastronomy.com/bad/movies/thecore_review.html

Zeff = Z – S = 12 – 10 = 2 Example 2: Magnesium Atom Z = 12 (the atomic number) S = approximately equal to the A neutral magnesium atom has 12 protons number of core electrons = 10 and 12 electrons Electronic configuration is 1s2 2s2 2p6 3s2 (Note: not dra wn to scale!)

12 protons and some neutrons (charge: +12) CORE All complete Net +2 inner shells of charge electrons ((gcharge: -10) 1s2 2s2 2p6 3s2 Outer electrons VALENCE beyond all SHELL complete inner Net -2 shells (charge: -2) charge Image modified from http://www.badastronomy.com/bad/movies/thecore_review.html

Clicker question #1

What is the effective nuclear charge (Zeff) of silicon, Si? (A) +1 (B) +2 (C) +3 (D) +4 (E) +5

Remove your name from mailing lists to reduce the amount of junk mail you receive. Core vs. Valence

An abbreviated periodic table (showing only the s- and p-blocks) H He

Zeff = 1-0 Zeff = 2-0

Li Be B C N O F Ne

Zeff = 3-2 Zeff = 4-2 Zeff = 5-2 Zeff = 6-2 Zeff = 7-2 Zeff = 8-2 Zeff = 9-2 Zeff = 10-2

Na Mg Al Si P S Cl Ar Zeff = 11-10 Zeff = 12-10 Zeff = 13-10 Zeff = 18-10

KCa GaGeAsSeBrKr

Eat less meat Recall Coulomb’s Law

Force of attraction (or repulsion): z Increases when magnitudes of charges increase z Decreases as distance between charges increases

Charge on proportionality positive part Charge on constant negative part k Q Q F = + − r2

Force of distance attraction between parts

To reason using Coulomb’s law, you must talk about the magnitudes of the9

charges (Q+ and Q-) and the separation of the charges (r).

Summary of Ionization Energy Trends

Ionization energy trends: • Increases from left to right across a period (row) • Decreases from top to bottom down a group (()column)

From Chemistry & Chemical Reactivity 5th edition by Kotz / Treichel. C 2003. Reprinted with permission of Brooks/Cole, a division of Thomson Learning: www.thomsonrights.com. Fax 800-730-2215.

Explaining trends in ionization energy

z Group trend z For the halogens, the ionization energy decreases as atomic number increases. z This is because… …the effective nuclear charge is the same (+7) for all the halogens, but the most loosely bound electron is in an energy level whose radius is larger as atomic number is larger, thus it is held less strongly due to greater r. z Period trend z For period 3, the ionization energy generally increases as atomic number increases. z This is because… …the most loosely bound electron is in the same energy level for all of them (more or less the same r away from the nucleus), but the effective nuclear charge increases as atomic number is larger, thus the most loosely bound electron is more tightly held as atomic number increases.

The next time a light bulb burns out, replace it with a fluorescent bulb. Chemical Explanations

In general, there are only a few basic concepts on which the logical steps of chemical explanations are built. The importance of size and charge (Coulomb’s law) 1. Core vs. valence in a single atom or ion The core is always positively charged and consists of all the protons plus the electrons that don’t participate in any action. All the electrons that participate in any action are in the valence shell. Comparisons are made based on magnitude of charges and distance separating the charges. (Note: it is possible to have competing effects.) 2. Charge density of an ion If two particles have equal charge but are different in size, the smaller one has greater charge density (more charge packed into a smaller space). Generally, something with greater charge density can have a stronger effect (e.g., it can get closer to oppositely charged particles, so the force of attraction will be greater) 3. Partial (polarizable) charge (next semester…) 12

Periods vs. Groups Comparing two elements in the same group: z Identify the trend in the property from top to bottom elements in the same group z Use # of shells argument to explain the trend z Different number of compp,()lete shells, so size (radius) of cores is different z Core charges are the same because valence electrons same z Arguments are usually based on distance between core and valence (r), or most loosely bound electron, being different while Q+ and Q- are the same Comparing two elements in the same period: z Identify the trend in the property from left to right elements in the same period

z Use Zeff arggpument to explain the trend z Same number of complete shells, so size (radius) of cores is the same z Different charges in nucleus, but same number of core electrons, leads to different core charge (Zeff) z Different numbers of electrons in valence

z Arguments are usually based on Q+ (core charge) and Q- (valence charge) being different while distance between core and valence (r) is nearly the same 13

Atomic Radius Comparisons

1. TddTrend down any g iven group?

2. Trend across any given period?

3. What h appens t o radii if these atoms become ions?

From Chemistry & Chemical Reactivity 5th edition by Kotz / Treichel. C 2003. Reprinted with permission of Brooks/Cole, a division of Thomson Learning: www.thomsonrights.com. Fax 800-730-2215.

Radii of Cations and Anions vs. Neutral Atoms

Gray = atoms Red = cations Blue = anions

Summary of Atomic Radii Trends

Atomic radius generally: z Increases down a group (column) z Decreases across a period (row)

Clicker question #2

Which answer is the correct completion of this statement?

As you go from left to right across a period in the periodic table, the atomic radius…

(A) decreases because the effective nuclear charge increases. (B) decreases because the effective nuclear charge decreases. (C) increases because the e ffec tive nuc lear c harge increases. (D) increases because the effective nuclear charge decreases.

Electron affinity The energy associated with adding an electron to an atom

z Electron affinities in kJ/mol z The more negative the electron affinity, the more exothermic. Therefore, the greater attraction the atom has for the electron to be added. z Trend: As you go from left to right across a row, the electron affinity generally ______

Electron affinity: two questions

z Why do He, Be, N and Ne have positive (endothermic) electron affinities? z Why is there not a steady trend in the electron affinities of the halogens? z F Cl Br I z -328, -349, -325, -295 kJ/mol z Look at the electron affinities in kJ/mol for the representative elements in the first five periods of the periodic table z The more negative the electron affinity, the greater the attraction of the atom for an electron z An electron affinity > 0 indicates that the negative ion is higher in energy than the separated atom and electron. z Electron affinity is the energy change that occurs when an electron is added to a gaseous atom z Energy is usually released when an atom gains an electron, giving the electron affinity a negative sign z Electron affinities generally become more negative from left to right across a row z Halogens obtain a noble gas structure when they gain an electron so their electron affinities are the most negative z The increase in electron affinity observed for group 5A elements relative to group 4A elements is due to the addition of an electron to a half-filled subshell.

The big picture: The pattern of a Coulomb’s law argument

1. Usually comparing one set of circumstances to a second set, to explain why one measure is larger or smaller than another z Neon atom vs. sodium atom with atomic radius

z MgCl2 vs. CaCl2 with energy required to break the ionic bonds 2. For each set of circumstances, determine what the relevant attraction is between a Q- and a Q+ z Attraction between outermost electron (-) and effective core charge (+) will affect atomic radius z Attraction between negatively charged ion (Cl -) and positively charged ion (Mg2+ or Ca2+) will determine strength of ionic bond 3. For each set of circumstances, determine what the distance of separation is between the + and – charges z Number of shells (periods) z Number of shells on + ion plus number of shells on – ion

4. Usually one variable, distance or (Q+ and Q-), can be considered constant while the other one varies. The one that varies is responsible for the difference in the measure z Neon has Q+=+8 while sodium has Q+=+1. Neon has 2 shells while sodium has 3 shells. Both differences lead to sodium’s outermost electron being further away and less tightly bound. z Both attractions are a +2 ion with a -1 ion. Cl- ion has same radius, but Mg2+ ion is smaller than Ca2+ ion, so separation distance between Q+ and Q- is smaller in MgCl2, therefore harder to break the ionic bond.

Electron Affinity The energy associated with an atom gaining an electron → Measures ease with which an atom gains an electron

H1s1 Trends in electron affinity He 1s2 Electron Affinities Li 1s2 2s1

2 2 H He Li Be Be C N O F Ne Na M Al Si P Si Cl Ar K Ca Ga Ge As Se Br Kr Rb Sr In Sn Sb Te In Xe Be 1s 2s 0 2 2 1 B1s2s 2p -50 2 2 2 C1s2s 2p -100 2 2 3 N1s2s 2p -150 l

2 2 4 o O 1s 2s 2p -200

2 2 5 kJ/m F1s2s 2p -250 2 2 6 Ne 1s 2s 2p -300 -350 Why don’t He, Be, N -400 and Ne gain electrons easily? Abbreviated Periodic Table

Electron Affinity

Two competing effects:

z Attraction between newly added electron and core Trend in electron affinity down a group thihititttdto which it is attracted Electron Affinities depends on distance between them –

as distance H He Li Be Be C N O F Ne Na M Al Si P Si Cl Ar K Ca Ga Ge As Se Br Kr Rb Sr In Sn Sb Te In Xe 0 increases, attraction -50 decreases -100 -150 z Repulsion between l o newly added electron -200

and other electrons kJ/m -250 already present – when electrons are -300 closer together (as -350 in smaller shells) -400 they repel each other more Abbreviated Periodic Table

Spend a few more hours every week doing something that requires no electricity or gas. Map of Chapter 8

z What holds ions together z Predicting qualitative trends z What holds molecules together z Predicting enthalpy of reaction from bond energies z Ionic vs. covalent character of bonds: polarity and electronegativity model z Lewis structure model z Simple structures (octet rule), with single and multiple bonds z Resonance structures z More complicated structures (breaking the octet rule) z Formal charges z Bond strength and length z Using Lewis structures to predict z Using Hess’s law and bond enthalpies

Ionic vs. Molecular (what you already know)

1. Ionic bonding z Occurs between ions z Ions are particles that are charged z Atoms that have gained or lost electron(s) z Polyatomic ions z In solid state, ions are arranged in regular, repeating, alternating + and - lattice structure z Strong attractions between + and -, and lattice structure in which every ion surrounded by many ions of opposite charge, make it difficult (energetically) to change solid to liquid z As a result, all ionically bonded materials are solids at room temperature and have very high melting points

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Ionic vs. Molecular (review)

2. Covalent bonding z Occurs between neutral atoms within individual molecules (neutral molecules are not charged) z Forces of attraction that hold atoms together inside of molecules result from each atom’s nucleus (+) attracting neighboring atom’s electrons (-) as well as its own electrons (also -) z Only valence electrons of atoms interact with valence electrons of other atoms z Covalent means ‘shared valence’ electrons z When covalently bonded molecules are in solid state, the molecules are arranged in regular, repeating lattice structures. The forces of attraction that hold together molecules in the solid state are weaker than ionic (full + and – monopoles) z inter-molecular forces are between polarized molecules (dipoles) or polarizable molecules ⇒ molecular solids have lower melting points than ionic solids

Lattice energy in ionic compounds is formation energy

Equal and opposite to energy required to break apart the lattice into separated ions

Unplug power transformers (chargers) when they are not in use. Period and Group Trends in Lattice Energy

Lattice Energies of Some Ionic Compounds

1200 Li Na 1000 K

800

600

of latticeenergy (kJ/mol) 400

200

Magnitude Magnitude 0 FClBrI Anion

Data from textbook p. 305

Qualitative trends (comparisons) in strength of ionic bonding

z Recall that Coulomb’s law predicts that force of attraction between two oppositely charged objects depends on magnitudes of()f charges (direct) and on distance that separates them (inverse squared)

z Comparing the lattice energies of two ionic compounds depends on two factors:

z Compare magnitudes of ionic charges (Q+ and Q-) over the same separation distance (r)

z Ctidit(Compare separation distance (r)ifth) if the same i on ic c harges (Q+ and Q-) z (Both factors can work in the same direction) z (If factors work in opposing directions, you need to know more quantitative information to make a prediction)

Bond polarity

z We will get to a model called the “electronegativity model,” which will helppp to explain wh y some molecular com pounds have more attractive forces between their molecules than other molecular compounds do z But first, it’s easier to talk about bonds in molecular compounds when we can say something about what attractions those bonds are made of z Remember, bonds in molecular compounds are called covalt(lent (co=sh are, valtlent=madflde of valence e ltlectrons ) z Generally, covalent bonds occur between nonmetals

Reduce standby electricity trickles. Plug seldom used appliances into a power strip and keep the power strip off except when you’re using an appliance.

Lewis Dot Structure Model

Valence number of an atom is the quantity of valence electrons z Periods 1, 2, 3 have only s and p electrons in valence shell z Period 3 and higher: d-orbitals are available and can play role in valence shell z Count only the s and p electrons in the valence shell to make Lewis dot structures Convention dictates “Hund’s rule” applies to filling the four locations that dots can occupy around a Lewis dot structure for an atom

(This is the Zeff model in different clothing/vocabulary)

H He +1C, -1V +2C, -2V Li Be B C N O F Ne +1C, -1V +2C, -2V +3C, -3V +4C, -4V +5C, -5V +6C, -6V +7C, -7V +8C, -8V Na Mg Al Si P S Cl Ar +1C, -1V +2C, -2V +3C, -3V +8C, -8V

As it warms up this summer, use air conditioning less often and open your windows to a breeze instead. Lewis Dot Structure Model

z Powerful model for predicting molecular shapes

z Assists in answering questions:

z Why do bonds between atoms within a molecule have specific angles?

z Why are molecular geometries complicated (3-dimensional) and not flat (2-dimensional)?

z Why are attractive forces between some molecules stronger/weaker than between other molecules?

z How is structure related to lots of other chemical and physical ppproperties of covalentl y bonded materials?

z Theory behind Lewis dot structures is that valence electrons are distributed as either

z Pairs of electrons that are shared by two atoms (shared pairs)

z Pairs of electrons that belong to a single atom (unshared, or lone pairs)

Lewis Structures of Simplest Molecules

F + F → F F

or There is no difference between one electron and anothWhher. When a mo lecu le is formed, the electrons F F have no allegiance to their original atoms. Lone pair of electrons Shared or bonding pair of electrons

Recycle more of your glass, metal cans, plastic containers, cardboard, newspaper, paper, etc. Octet Rule

A noble gas configuration (8 valence electrons) is stable. Atoms tend to form molecules or polyatomic ions in such a way that each atom is surrounddb8ded by 8 e lec trons (an oc tt)tet). ¾ There are lots of exceptions. They are interesting and we will study them. But first let’s master the simplest rules.

O + O → OO OOO

Conservation of electrons: Double How many electrons on left? bond How many electrons on right?

Lewis Dot Structure Model

z Theory of Lewis dot structures is that valence electrons are distributed as either

z Pairs of electrons that are shared by two atoms (shared pairs)

z Pairs of electrons that belong to a single atom (unshared, or lone pairs) Lone pair of electrons (belongs exclusively to this F F oxygen) O O

Lone pair of electrons Shared or Double bond (shared bonding pair of electrons get counted as electrons belonging to both oxygen atoms)

Think – Pair – Share

Draw the Lewis structures of the following molecules. Don’t forget to show all the valence electrons: both the bonding electrons and the lone pairs. (It might help to check your work by conservation of valence electrons.)

1) N2

2) CH4

3) CH3Cl

4) NH3

5) H2CO

Reminder: Electron Affinity

Electron affinity measures how much an atom “likes” electrons

Except not the noble gases

Building Lewis Structures

1. Determine central atom (atom with lowest electron affinity because electron densityyp will spread as far as possible toward the extremities of a molecule, given the opportunity)

2. Count total number of valence electrons in molecule

3. Arrange atoms around central atom

4. Start with single bonds

5. Place remaining valence electrons

6. Move electrons to form octets, making double or triple bonds where necessary Check: Make sure you have conservation of electrons

Hang laundry to air dry instead of using a dryer.

Buy fresh local produce instead of industrially produced imported. Practicing the Steps

+ Draw the Lewis structure for ammonium (NH4 ) 1. N is central because H atoms can only bond once 2. Total valence electrons • 5 on N • 1 on each H gives 4 more • Ion has +1 charge, meaning one electron is removed • Total = 5 + 4 – 1 = 8 The rest of the steps: + H

H N H H

Practicing the Steps

– Draw the Lewis structure for the nitrate ion (NO3 ) 1. N is central because it has the least electron affinity 2. Total valence electrons • 5 on N • 6 on each of three O atoms gives 18 more • Ion has -1 charge, meaning one electron is added • Total = 5 + 18 + 1 = 24 The rest of the steps:

O N O O Take the T more often instead of driving

More Practice

Draw these Lewis structures + a) NO2 b) CN– c) SCN–

d) O3 Then answer the following two questions:

1. Which two of the above structures are isoelectronic? 2. What do you notice that is special about isoelectronic species?

You will need the following definition: Isoelectronic = same number of (valence) electrons and same number of atoms to distribute them around Switch to opaque projector

Predictions that can come from Lewis structures

z Resonance z Bond order (useful for comparing strengths of bonds) z Bond length comparisons (comparing one molecule to another) z Formal charge (useful for predicting which resonance structures are most stable)

Resonance Structures

– – – O N O O N O O N O

O O O

z What do these structures have OON in common? Average bond z How are they different? order = 4 bonds split over 3 O z Which of these is the actual locations = 4/3 - structure of NO3 ? Note: This is not a correct Lewis structure. It is # of bonds drawn this way only to Bond order = emphasize the bond # of locations order.

Bond Order and Bond Length/Strength

z Bond order z Single bond is bond order 1 O N O z Double bond is bond order 2 Average bond order = 4 bonds z Triple bond is bond order 3 split over 3 O z Fractional bond orders occur locations = 4/3 when there are resonance structures z Bond strength z The greater the bond order, the stronger the bond (the more energy required to break the bond) z Bond length z The greater the bond order, the shorter the bond length

Bond strength

Bond length

Warning: Two other models seem similar to the formal charges model but are completely different models. All three models are ways of understanding the Formal distribution of electron density to explain some chemical or physical properties, but they are all used for different purposes. Charges • Oxidation numbers model for understanding redox reactions • Electronegativity model for explaining polarity (ionic character) in covalent bonds

A comparison between the valence electrons originally contributed by an atom and the electrons that it looks like the atom would have if all bonds were broken and electrons reassigned democratically.

The valence Electrons assigned Therefore, the formal electrons that democratically if bonds charges on each atom were originally hypothetically broken are contributed: -1 7 4 6 – +1 – • each O had 6 0 O N O O N O • the N had 5

The formal charges O O model is used to 7 explain why some -1 resonance structures are more stable than Notice that the sum of the formal others. charges must equal the ion charge

Formal Charges and Alternative Structures

z If more than one Lewis structure exists, the most stable structure is the one in which the formal charges make most sense z Negative formal charges on atoms with large electron affinity z Positive formal charges on atoms with small ionization energies (small electron affinity)

– -1 – -1 0 O C N O C N 0 0 0

-2 – +1 0 OCN

What we have learned so far

z Ionic vs. covalent bonding z Ionic bond is attraction between + and – ions, held toggyether by Coulomb force of attraction acting across ion separation z Molecular bond is valence electrons shared between two atoms and attracted (but not necessarily equally) to both atoms in the bond z Many aspects of molecular bonding can be modeled by Lewis structures z Bond order/strength and length z Resonance structures of molecules (or ions) and formal charges on individual atoms in the molecules (or ions) z Bond polarity z Molecular geometry