Honors Chemistry Final

Chapter 5: Thermochemistry

5.1 The Nature of Energy

• Kinetic Energy: The energy of motion.

  • Mass and energy are directly proportional.

• Potential Energy: The energy an object possesses by virtue of its position.

  • Can be converted into kinetic energy.

  • Includes electrostatic potential energy ($E_{el}$), which is the most important form in molecules.

  • Governed by Coulomb's Law (see diagram on p. 168).

Units of Energy

• SI Unit: Joule (J)

  • $1 J = 1 kg [\frac{m^2}{s^2}]$

• Calorie (cal): Commonly used, but not SI.

  • $1 cal = 4.184 J$ (exactly)

• Nutritional Calorie (Cal): Equivalent to a kilocalorie (kcal).

  • $1 Cal = 1000 cal = 1 kcal$

Systems and Surroundings

• System: The part of the universe of interest.

  • Example: Reactants in a chemical reaction.

• Surroundings: The rest of the universe outside the system.

  • Example: Beaker containing reactants and everything beyond.

Transferring Energy: Work and Heat

• Work: Energy transferred by causing the motion of an object against a force.

  • Force: Push or pull on an object.

  • Work: Product of force applied to an object over a distance.

• Heat: Energy transferred between two objects due to a temperature difference.

• Energy: Capacity to do work or transfer heat.

5.2 The First Law of Thermodynamics

• Internal Energy (E): Total energy of a system.

  • Cannot be measured directly; changes in internal energy ($\Delta E$) are observed.

Relating ΔE to Heat and Work

• First Law of Thermodynamics: Energy cannot be created or destroyed; the total energy of the universe (system + surroundings) is constant.

  • $\Delta E = q + w$

    • q = heat added to or absorbed by the system

    • w = work done on or by the system

  • Energy deposited into system: \Delta E > 0

  • Energy withdrawn from system: \Delta E < 0

• Sign Conventions:

  • Heat gained: q > 0

  • Heat lost: q < 0

  • Work done on system: w > 0

  • Work done by system: w < 0

Exothermic and Endothermic Processes

• Endothermic: System absorbs heat from surroundings.

  • Feels cold.

• Exothermic: System transfers heat to surroundings.

  • Feels hot.

State Functions

• State Function: A property that depends only on the initial and final states of the system, not the path taken.

  • Internal energy ($\Delta E$) is a state function.

  • Example: Discharging a battery.

    • Shorting a battery with a coil (no work): All energy lost as heat (q < 0, w = 0).

    • Battery turns a fan (work is done): Heat and work are involved (q < 0, w < 0).

    • In both cases, $\Delta E$ is the same because it's a state function.

5.3 Enthalpy

• Enthalpy (H): Heat transferred between system and surroundings at constant pressure, with no forms of work other than pressure-volume changes.

  • $H = E + PV$

• State Function: Yes, enthalpy is a state function.

• Extensive Property: Magnitude of $\Delta H$ is directly proportional to amount.

• Constant Pressure: If a process occurs at constant pressure:

  • $\Delta H = q_p$

• Sign Conventions:

  • \Delta H > 0: System gains heat (endothermic).

  • \Delta H < 0: System loses heat (exothermic).

• Work and Enthalpy: At constant pressure:

  • $-w = P\Delta V$ (piston moves to maintain constant pressure)

5.4 Enthalpies of Reaction

• Thermochemical Equations: Balanced chemical equations with $\Delta H$ values.

• Enthalpy as an Extensive Property: Magnitude of $\Delta H$ is directly proportional to the amount of reactants consumed or products formed.

  • Example:

    • $CH<em>4(g) + 2O</em>2(g) \rightarrow CO<em>2(g) + 2H</em>2O(g) \quad \Delta H = -802 kJ$

    • $2CH<em>4(g) + 4O</em>2(g) \rightarrow 2CO<em>2(g) + 4H</em>2O(g) \quad \Delta H = -1604 kJ$

• Reversing a Reaction: Changes the sign of $\Delta H$.

  • Example:

    • $CO<em>2(g) + 2H</em>2O(l) \rightarrow CH<em>4(g) + 2O</em>2(g) \quad \Delta H = +890 kJ$

    • $2O<em>2(g) + CH</em>4(g) \rightarrow CO<em>2(g) + H</em>2O(l) \quad \Delta H = -890 kJ$

• Enthalpy Depends on State: Change in enthalpy depends on the physical states of reactants and products.

  • Example:

    • $2O<em>2(g) + CH</em>4(g) \rightarrow CO<em>2(g) + H</em>2O(l) \quad \Delta H = -890 kJ$

    • $2O<em>2(g) + CH</em>4(g) \rightarrow CO<em>2(g) + H</em>2O(g) \quad \Delta H = -802 kJ$

    • $H<em>2O(l) \rightarrow H</em>2O(g) \quad \Delta H = +88 kJ$

    • Less heat is available for transfer to the surroundings because the enthalpy of $H<em>2O(g)$ is greater than that of $H</em>2O(l)$.

Enthalpy as a Guide

• Spontaneous Process: Thermodynamically favored to happen.

  • A reaction with a large, negative $\Delta H$ will be spontaneous. (exothermic)

  • A reaction with a small, positive $\Delta H$ will be spontaneous in reverse. (endothermic)

• Note: The sign of $\Delta H$ indicates the direction of energy flow, not a positive or negative value for energy.

5.5 Calorimetry

• Calorimetry: Measurement of heat flow.

• Calorimeter: Apparatus that measures heat flow.

• Heat Capacity (C): Amount of energy required to raise the temperature of an object by one degree.

• Molar Heat Capacity ($C_m$): Heat capacity of 1 mole of a substance.

• Specific Heat ($C_s$): Specific heat capacity; heat capacity of 1 gram of a substance.

  • $q = m \times C_s \times \Delta T$

Constant Pressure Calorimetry

• Atmospheric pressure is constant.

• Fun Fact: $1 cal = 4.184 J$

• Bomb Calorimeter: Used to measure the heat of combustion of a particular reaction at constant volume.

  • Bomb calorimeters have to withstand the large pressure within the calorimeter as the reaction is being measured.

5.6 Hess's Law

• Hess's Law: If a reaction is carried out in a number of steps, $\Delta H$ for the overall reaction is the sum of $\Delta H$ for each individual step.

  • The quantity of heat generated is independent of the number of steps needed for the reaction to occur.

  • Example:

    • $CH<em>4(g) + 2O</em>2(g) \rightarrow CO<em>2(g) + 2H</em>2O(g) \quad \Delta H = -802 kJ$ (Combustion reaction)

    • $2H<em>2O(g) \rightarrow 2H</em>2O(l) \quad \Delta H = -88 kJ$ (Condensation of water)

    • $CH<em>4(g) + 2O</em>2(g) \rightarrow CO<em>2(g) + 2H</em>2O(l) \quad \Delta H = -890 kJ$ (Overall reaction)

    • $\Delta H<em>1 = \Delta H</em>2 + \Delta H_3$

5.7 Enthalpies of Formation

• Enthalpy of Formation ($\Delta H_f$): Enthalpy change that occurs if 1 mol of a compound is formed from its constituent elements.

• Standard Conditions: 1 atm and 25 °C (298 K).

• Standard Enthalpy ($\Delta H^o$): Enthalpy measured when everything is in its standard state.

• Standard Enthalpy of Formation ($\Delta H^o_f$): Enthalpy change when 1 mol of compound is formed from substances in their standard states.

• Standard States: If there is more than one state for a substance under standard conditions, the more stable one is used.

• $\Delta H^o_f$ of the most stable form of an element is zero.

• Diatomics exist, refer to periodic table info from chapter 1.

We use Hess’ Law to calculate enthalpies of a reaction from enthalpies of formation:

$\Delta H<em>{rxn}^o = \Sigma n \Delta H</em>f^o(products) - \Sigma m \Delta H_f^o(reactants)$

• where n and m are stoichiometric coefficients

5.8 Bond Enthalpies

Bond Enthalpy: The enthalpy associated with braking one mole of a particular bond in a gaseous substance.

• The bond enthaply is always positive because energy is required to break chemical bonds.

• Energy is always released when a bond forms between gaseous fragments.

• The greater the bond enthaply, the stronger the bond.

*bond enthalpies

$\Delta H_{rxn} = \Sigma (bond \, enthalpies \, of \, bonds \, broken) - \Sigma (bond \, enthalpies \, of \, bonds \, formed)$

Chapter 6: Electronic Structure of Atoms

6.1: The Wave Nature of Light

• Electromagnetic Radiation (EMR): Carries energy through space.

• Speed of EMR: Constant in a vacuum: $3.00 \times 10^8 m/s$.

• Wave-like Characteristics:

  • Wavelength ($\lambda$): Distance from crest to crest.

  • Amplitude (A): Height of the wave.

  • Frequency ($\nu$): Number of cycles passing a point per second.

• The wave nature is due to periodic oscillations of the intensitites of electronic and magnetic forces associated with the radiation.

• Frequency and wavelength are inversely related.

Speed of a Wave

• The speed is given by frequency multiplied by wavelength:

  • $c = \nu \lambda$

  • For light, $c = 3.00 \times 10^8 m/s$

The Electromagnetic Spectrum

• Interaction of radiation with matter led to modern atomic theory

*Differences in absorption or emission of photons in different spectral regions are related to the different types of molecular motion or electronic transition:

*Microwave radiation is associated with transitions in molecular rotational levels.

*Infrared radiation is associated with transitions in molecular vibrational levels.

*Ultraviolet/visible radiation is associated with transitions in electronic energy levels.

6.2: Quantized Energy and Photons

• Problems with Wave Theroy:

  • Emission of light from hot objects.
    Black body radiation: stove burner, light bulb filament.

  • The photoelectric effect.

  • Emission spectra

• Planck: Energy can only be absorbed or released from atoms in certain amounts called quanta.

  • $E = h\nu$

    • h is Planck's constant ($6.626 \times 10^{-34} J \cdot s$)

    • Quantization Analogy: Walking up a ramp vs. stairs shows continuous vs. quantized change in height.

The Photoelectric Effect

• Evidence for the particle nature of light (quantization).

• Light shines on a metal surface -> electrons are ejected.

• A minimum threshold frequency must be reached.

  • Below this, no electrons are ejected.

  • Above this, the # of electrons ejected depends on the intensity of the light.

*Einstein assumed that light traveled in energy packets called photons

• The energy of one photon:
  *Energy and frequency are directly proportional
  *Therefore radiant energy must be quantized!

6.3: Line Spectra and the Bohr Model

• Line Spectra: Radiation that spans an array of different wavelengths is continuous.

  • White light through a prism creates a continuous spectrum of colors.

  • Spectra tube emits light unique to the element in it.

    • Looking at it through a prism, only lines of a few wavelengths are seen.

    • Black regions correspond to wavelengths that are absent.

Bohr Model

• Rutherford model: electrons orbit the nucleus like planets around the sun.

• Physics: a charged particle moving in a circular path should lose energy; therefore, the atom should be unstable.

• 3 Postulates for Bohr's theory:

  1. Electrons move in orbits that have defined energies.

  2. An electron in an orbit has a specific energy.

  3. Energy is only emitted or absorbed by an electron as it changes from one allowed energy state to another ($E = h\nu$).

*Since the energy states are quantized, the light emitted from excited atoms must be quantized and appear as line spectrum
*Bohr Showed that $$E=-Rh(\frac{1}{n^2})$$
n = energy level

• The first orbit in the Bohr model has n = 1, is closest to the nucleus, and has the lowest energy (ground state)

• The furthest orbit in the Bohr model has n close to infinity

• The first orbit in the Bohr model has n = 1, is closest to the nucleus, and has the lowest energy (ground state).

• The furthest orbit in the Bohr model has n close to infinity.

• Electrons in the Bohr model can only move between orbits by absorbing and emitting energy in quanta ($h\nu$).

• The amount of energy absorbed or emitted on movement between states is given by $S = \frac{-Rh}{n^2}$

Line Spectra

• The Balmer series for Hydrogen ($n_f = 2$, which is in the visible region).

• The Rydberg Equation (P. 220) allows for the calculation of wavelengths for all the spectral lines $(\frac{1}{\lambda} = R<em>H(\frac{1}{n</em>1^2}-\frac{1}{n_2^2})$ n = energy level

• $n_f$ values for other regions of the EMS:

  • Lyman (UV): $n_f = 1$

  • Paschen (IR): $n_f = 3$

  • Brackett (IR): $n_f = 4$

  • Pfund (IR): $n_f = 5$

• Note the units

Limitations of the Bohr Model

• It can only explain the line spectrum of hydrogen adequately.

• Electrons are not completely described as small particles.

• It doesn't account for the wave properties of electrons.

6.4:The Wave Behavior of Matter

• Using Einstein's and Planck's equations, de Broglie showed:

  • $\lambda = \frac{h}{mv}$

• The momentum, mv, is a particle property, but $\lambda$ is a wave property.

6.5: Quantum Mechanics and Atomic Orbitals

• Heisenberg's Uncertainty Principle: We cannot determine exactly the position, direction of motion, and momentum of an electron simultaneously.

• Erwin Schrödinger proposed an equation that contains both wave and particle terms.

  • Solving the equation leads to wave functions (shape of the electronic orbital).

Orbitals and Quantum Numbers

• If we solve the Schrödinger equation, we get wave functions and energies for the wave functions.

• We call wave functions orbitals (regions of highly probable electron location).

• Schrödinger's equation requires 3 quantum numbers:

  1. Principal Quantum Number (n): This is the same as Bohr's n. As n becomes larger, the atom becomes larger and the electron is further from the nucleus.

  2. Azimuthal Quantum Number (l): This quantum number depends on the value of n. The values of l begin at 0 and increase to (n - 1). We usually use letters for l (s, p, d, and f for l = 0, 1, 2, and 3). Usually, we refer to the s, p, d, and f-orbitals (the shape of the orbital) (AKA “subsidiary quantum number”).

  3. Magnetic Quantum Number ($m_l$): This quantum number depends on l: The magnetic quantum number has integral values between -l and +l. Magnetic quantum numbers give the 3D orientation of each orbital.

6.6: Representations of Orbitals

The s-Orbitals: All s-orbitals are spherical
As n increases, the s-orbitals get larger

The p-orbitals: There are three p-orbitals px, py, and pz
The three p-orbitals lie along the x-, y- and z- axes of a Cartesian system The letters correspond to allowed values of ml of -1, 0, and +1 The orbitals are dumbbell shaped
As n increases, the p-orbitals get larger
The d and f orbitals: There are five d and seven f-orbitals

Max = (n-1)
l → letter (-l to +l)
(n2) (# of orientations)

6.7: Many-Electron Atoms

• Orbitals of the same energy are said to be degenerate.

• For n ≥ 2, the s- and p-orbitals are no longer degenerate because the electrons interact with each other.

• Therefore, the Aufbau diagram looks slightly different for many-electron systems.

Electron Spin and the Pauli Exclusion Principle
Since electron spin is also quantized, we define
ms = spin quantum # = ± ½

Pauli’s Exclusion Principle: no two electrons can have the same set of 4 quantum numbers Therefore, two electrons in the same orbital must have opposite spins Electron capacity of sublevel = 4l + 2 Electron capacity of energy level = 2n2

6.8: Electron Configurations Hund’s Rule

• Ground state electron configurations tell us in which orbitals the electrons for an element are located.

• Lowest energy state

• Three rules are applied:

  • Aufbau Principle

  • Pauli’s Exclusion Principle

  • Hund's Rule: for degenerate orbitals, electrons fill each orbital singly before any orbital gets a second electron.

• A paramagnetic atom has one or more unpaired electrons.

• A diamagnetic atom is on in which all electrons are paired.

• Paramagnets do not retain magnetization in the absence of a magnetic field, because thermal energy randomizes electron spin orientations.

• Diamagnetic atoms repel magnetic fields. The unpaired electrons of paramagnetic atoms realign in response to external magnetic fields and are therefore attracted.

6.9: Electron Configurations and the Periodic Table

• The periodic table can be used as a guide for electron configurations.

• The period number is the value of n.

• The periodic table can be used as a guide for electron configurations.

• The period number is the value of

• The periodic table can be used as a guide for electron configurations.

• The period number is the value of n.

of columns is related to the number of electrons that can fit in the subshells

• Inner shell or core electrons

Electron State POGIL

The following snips from a POGIL activity will help you to understand the excited state (promoted electrons) and ground state (the regular, expected electron configuration).

Chapter 7: Periodic Properties of the Elements

• Dimitri Mendeleev and Lothar Meyer arranged the elements in order of increasing atomic weight.

• Modern periodic table: arrange elements in order of increasing atomic number.

  • As properly belonged underneath P and not Si

  • A missing element underneath Si

  • He predicted a number of properties for this element

  • In 1886 Ge was discovered

  • Properties of Ge match Mendeleev’s predictions

7.2: Effective Nuclear Charge

• Coulombic Force: Force of attraction between two charged particles.

  • Increases with increasing charge and/or decreasing distance.

• Effective Nuclear Charge ($Z_{eff}$): Net positive charge experienced by an electron in a many-electron atom.

  • Not the same as the charge on the nucleus due to the effect of inner electrons.

  • Electrons are attracted to the nucleus but repelled by the electrons that screen them from the nuclear charge

  • The nuclear charge experienced by an electron depends on
    (1) its distance to the nucleus
    (2) the number of core electrons

$Z_{eff} = Z - S$
Z is # of protons or atomic #
S is the amount of shielding

• Valence electrons experience less than the full nuclear charge (shielding effect).
  $Z<em>{eff}$ increases across period. Core electrons are the same/constant. The added electron does not impact shielding the other electrons from the positive pull of the nucleus. $Z</em>{eff}$ is fairly constant down a group
  Valence electrons are constant so number of core electrons are increasing

1. 3: Sizes of Atoms and Ions
   Distance between two nuclei in a diatomic molecule is called the bond distance
   Half of the bond distance is called the covalent radius of the atom

Periodic Trends in Atomic Radi

• The properties of elements vary periodically.

• Atomic radius increases down a group (more noticeable).

• Atomic radius decreases across a period.

• There are two factors at work:

  • Principal quantum number

  • Electrons fill higher energy levels

  • Effective nuclear charge

  • Electrons fill the same energy level, no increase in shielding

For ions of the same charge, ion size increases down a group.
Electrons are in higher energy levels.
All the members of an isoelectronic series have the same number of electrons (O2- , F- , Na+, Mg2+, Al3+)
As nuclear charge increases in an isoelectronic series, the ions become smaller: O2- > F- > Na+ > Mg2+ > Al3+
Added electrons are in the same energy level, so no increase in shielding
Removed electrons result in an increasingly positive nucleus which pulls electrons toward it Problem Solving:

7.4: Ionization Energy

• The first ionization energy, I1, is the amount of energy required to remove an electron from a gaseous atom:

  • $Na(g) \rightarrow Na^+(g) + e^-$

• The second ionization energy, I2, is the energy required to remove a second electron from a gaseous ion:

• $Na^+(g) \rightarrow Na^{2+}(g) + e^-$

• The larger the ionization energy, the more difficult it is to remove the electron

Electron Configuration of Ions

• Cations: Electrons removed from the orbital with the highest principle quantum number, n, first:
  * $Li (1s^2 2s^1) \Rightarrow Li^+ (1s^2)$
  * $Fe ([Ar]3d^6 4s^2) \Rightarrow Fe^{3+} ([Ar]3d^5)$

• Anions: Electrons added to the orbital with the highest n:
  * $F (1s^2 2s^2 2p^5) \Rightarrow F^- (1s^2 2s^2 2p^6)$

7.5: Electron Affinities

• Think: does it want to gain an electron??

• Electron affinity is the energy change when a gaseous atom gains an electron to form a gaseous ion:
  * $Cl(g) + e^- \rightarrow Cl^-(g)$

• Electron affinity can either be exothermic (above) or endothermic:
  * $Ar(g) + e^- \rightarrow Ar^-(g)$

EA usually becomes MORE negative across a period.
Greater attraction between the atom and the added electron Groups 2, 15 and 18 EA has no real group trend
Increasing Metallic Character

7.6: Metals, Nonmetals, and Metalloids

Metals Metallic character refers to the properties of metals:
Shiny or lustrous Malleable and ductile Oxides form basic ionic solids Form cations in aqueous solution
Metallic character increases down a group Metallic character decreases across a period
Metals have low ionization energies. Most neutral metals are oxidized rather than reduced Form cations
Most metal oxides are basic:
Metal oxide + water → metal hydroxide
The greater the metallic character of the metal, the stronger the basic character of the oxide

Nonmetals: Nonmetals are more diverse in their behavior than metals Properties are opposite those of metals Compounds composed only of nonmetals are molecular When nonmetals react with metals, nonmetals tend to gain electrons metal + nonmetal → salt Most nonmetal oxides are acidic:
nonmetal oxide + water → acid
Nonmetal oxide + base → salt + water The greater the nonmetallic character of the central atom, the stronger the acidic character of the oxide Metalloids Metalloids have properties that are intermediate between metals and nonmetals Si has a metallic luster but it is brittle Semiconductors Problem Solving:

7.7: Group Trends for the Active Metals

Group 1A: The Alkali Metals Alkali metals are all soft Chemistry is dominated by the loss of their single s electron M → M+ + e-
Reactivity increases down the group because metallic character increases down a group Alkali metals react with water to form MOH and hydrogen gas
2M(s) + 2H2O(l) → 2MOH(aq) + H2(g) Based on reactivity, RbOH would be stronger than KOH and KOH would be stronger than NaOH
Flame Test:
The s electron is excited by the flame and emits energy when it returns to the ground state.

Group 2A: The Alkaline Earth Metals Alkaline earth metals are harder and more dense than the alkali metals
The chemistry is dominated by the loss of two s electrons M → M2+ + 2e- Mg(s) + Cl2(g) → MgCl2(s) Increasing activity down the group Be does not react with water
Mg will only react with steam
Ca, Sr and Ba: M(s) + 2H2O(l) → M(OH)2(aq) + H2(g

7. 8: Group Trends for Selected Nonmeta
   Hydrogen Colorless diatomic gas, H2 High IE because there are no inner electrons It can either gain another electron to form the hydride ion, H− , or lose its electron to become H+:
   2Na(s) + H2(g) → 2NaH(s) 2H2(g) + O2(g) → 2H2O(g) H+ is a proton
   The aqueous chemistry of hydrogen is dominated by H+(aq)

Group 16: The Oxygen Group Metallic character increases down a group (O2 is a gas, Te is a metalloid, Po is a metal) Two important forms of oxygen: O2 and ozone, O3 Ozone can be prepared from oxygen:
3O2(g) → 2O3(g) ΔH = +284.6 kJ. Ozone is pungent and toxic
Oxygen Group:
Most common form is yellow S8.
Sulfur tends to form S2- in compounds.
Allotropes: Different forms of the same element in the same state.

The chemistry of the halogens is dominated by gaining an electron to form an anion: X2 + 2e- → 2X- Fluorine is one of the most reactive substances known All halogens consists of diatomic molecules, X2

Group 18: The Noble Gases These are all nonmetals and monatomic Unreactive because they have completely filled s and p subshells In 1962 the first compound of the noble gases was prepared: XeF2, XeF4, and XeF6 Xe has a low enough IE to react with substances that readily remove electrons To date, the only other noble gas compounds known are KrF2 and HArF

Chapter 8 Basic Concepts of Chemical Bonding

8.1: Chemical Bonds, Lewis Symbols, and the Octet Rule

• Chemical Bond: Attractive force holding two or more atoms together.

• Covalent Bond: Electrons are shared.

  • Usually found between nonmetals.

• Ionic Bond: Results from the transfer of electrons from a metal to a nonmetal.

• Metallic Bond: Attractive force holding pure metals together.
  Lewis Dots
  Valence Electrons are represented as dots around the symbol for the element Electrons available for bonding are indicated by unpaired dots
  All noble gases except He have an ns2np6 configuration Octet rule: atoms tend to gain, lose, or share electrons until they are surrounded by 8 valence electrons (4 electron pairs)

8.2: Ionic Bonding

Lattice energy increases (and thus stability increases) with:
(1) charge on the ions (2) decreasing distances between the ions
Trend: Ionic radius increases down a group of the periodic table
Makes ionic compounds hard, high melting

• NaCl forms a very regular structure, Regular arrangement of Na+ and Cl- in 3D, Ions are packed as closely as possible

Energetics of Ionic Bond Formation

• Lattice Energy: The energy required to completely separate an ionic solid into its gaseous ions.

  • $NaCl(s) \rightarrow Na^+(g) + Cl^-(g)$

    • This process is endothermic ($\Delta H = +788 kJ/mol$).

• Lattice energy depends on the charges on the ions and the sizes of the ions.

Lattice Energy Increases.. Charges on the ions increase Ionic distance between the ions decreases (think Coulombs law) Because ionic radius increases down a group, lattice energy DECREASES down a group. High lattice energies = hard, brittle, high melting point.

8. Born-Haber Cycle Many factors affect the energy of ionic bonding.

Starts with the metal and nonmetal elements Na(s) and Cl2(g) then makes gaseous atoms
Then makes ions and combine them into NaCl(s).
Born–Haber cycles are used primarily as a means of calculating lattice energy which cannot otherwise be measured directly

8.3: Covalent Bonding

Electronegativity is a sliding scale that determines polarity. Similar atoms=covalent bonded neither want to gain or lose e-. Unequal sharing of e- creates polar bonds or ionic bonds
Sharing of electrons to form a covalent bond does not imply equal sharing of those electrons
Unequal sharing of electrons results in polar bonds

Lewis Structures Covalent bonds can be represented by the Lewis symbols of the elements
Pairs in a bond are represented by a single line
Multiple bonds are possible!
More than one pair of electrons is shared between two atoms Single, Double, Triple bonds. Bond distances decrease from Single to triple bonds

8.4: Bond Polarity and Electronegativity

Sharing of electrons to form a covalent bond does not imply equal sharing of those electrons
Unequal sharing of electrons results in polar bonds Nonpolar Covalent Bonds Polar Covalent Bonds lonic Bonds
Electronegativity: The ability of one atom in a molecule to attract electrons to itself Pauling scale is from 0.7 (Cs) to 4.0 (F) Electronegativity increases across a period Electronegativity decreases down a group

Molecule contains polar bonds, but it is not itself polar; the average location of positive/negative charges are the same

8.5: Drawing Lewis Structures

Polyatomic cations: Take away number of electrons equal the positive charge
Polyatomic anions: Add number of electrons equal to the negative charge

Formal Charge: FC = valence electrons - # lone pair electrons - # bonding e-/2 Most stable structures.. All atoms have lowest formal charge Negative formal charges assigned to most electronegative

8.6: Resonance Structures

Draw more than one structure with the octet rule being obeyed to determine which structure is most resonable

Exceptions to the Octet Rule Odd Number of Electrons Few examples
Dimerizations
Less than an Octet Relatively rare Compounds of Groups 2 and 13 Greater than an Octet This is the largest class of exceptions Atoms from the 3rd period onward can accommodate more than an octet If everything has an octet and electrons remain, put them on the central atom Strengths of Covalent Bonds The energy required to break a covalent bond is called the bond dissociation enthalpy, D
Bond enthalpies can either be positive or negative

Strengths of Covalent Bonds:

the energy required to break a covalent bond is
Is called the Bond-Dissociation Enthalpy.

• Δh has to be positive.

• In any chemical reaction, bonds break THEN form, which relates to delta h:

Bonds breaking/forming
P-P : 200 kl