Chapter 11-study guide PDF

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Notes for the third partial.
It was part of the final exam...

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11. 1 Lewis Structure of Elements and Compounds Bonding Theories • Predict how atoms bond together to form compounds. • Predict what combinations of atoms form compounds and what combinations do not. • Predict why salt is NaCl and not NaCl2 and why water is H2O and not H3O. • Explain the shapes of molecules, which in turn determine many of their physical and chemical properties. Bonding Theory: Lewis Theory  Named after G. N. Lewis (1875–1946), the American chemist who developed it.  Represent electrons as dots and draw “dot structures”, or Lewis structures, to represent molecules.  These structures have tremendous predictive power.  Lewis theory can determine whether aparticular set of atoms will form a stable molecule and what that molecule might look like.  Although modern chemists also use more advanced bonding theories to better predict molecular properties, Lewis theory remains the simplest method for making quick, everyday predictions about molecules. Representing Valence Electrons with Dots • In Lewis theory, the valence electrons of main-group elements are represented as dots surrounding the symbol of the element.  The result is called a Lewis structure, or dot structure. Writing Lewis Structures for Elements

Lewis Structures of Ionic Compounds: Electrons Transferred:

Covalent Lewis Structures: Covalent Compounds

In water, hydrogen and oxygen share their electrons so that each hydrogen atom gets a duet and the oxygen atom gets an octet. Covalent Lewis Structures: Electrons Shared The shared electrons—those that appear in the space between the two atoms—count toward the octets (or duets) of both ofthe atoms.

Lewis Theory Explains Why the Halogens Form Diatomic Molecules

Double Bonds: In Lewis Theory, Atoms Can Share More Than One Electron Pair to Attain an Octet Each oxygen atom now has an octet because the additional bonding pair counts toward the octet of both oxygen atoms. Triple Bonds: In Lewis Theory, Atoms Can Share More than One Electron Pair to Attain an Octet. Writing Lewis Structures for Covalent Compounds 1. Write the correct skeletal structure for the molecule (continued). Remember two guidelines.  First,hydrogen atoms will always be terminal. Since hydrogen requires only a duet, it will never be a central atom because central atoms must be able to form at least two bonds and hydrogen can form only one.  Second,many molecules tend to be symmetrical, so when a molecule contains several atoms of the same type, these tend to be in terminal positions.  This second guideline has many exceptions. In cases where the skeletal structure is unclear, the correct skeletal structure is usually provided. 2. Calculate the total number of electrons for the Lewis structure by summing the valence electrons of each atom in the molecule.  The number of valence electrons for any main-group element is equal to its group number in the periodic table.  If you are writing a Lewis structure for a polyatomic ion, the charge of the ion must be considered when calculating the total number of electrons. 3. Distribute the electrons among the atoms, giving octets (or duets for hydrogen) to as many atoms as possible.  Begin by placing two electrons between each pair of atoms. These are the minimal number of bonding electrons.  Then distribute the remaining electrons, first to terminal atoms and then to the central atom, giving octets to as many atoms as possible.  If any atoms lack an octet, form double or triple bonds as necessary to give them octets. Do this by moving lone electron pairs from terminal atoms into the bonding region with the central atom. Formal Charges (FC) 1. Formal charge (FC) is a fictitious charge assigned to an atom in a Lewis structure, assuming that electrons in all chemical bonds are shared equally between atoms. This helps to determine the best Lewis structure. 2. During bonding, atoms may end with more or fewer electrons than the valence electrons they brought in order to fulfill octets. This results in atoms having a formal charge.

FC= valence e- -nonbonding e- -1/2 bonding eWrite the Lewis Structure for the CN− Ion Write the skeletal structure: CN

Calculate the total number of electrons for the Lewis structure by summing the number of valence electrons for each atom and adding one for the negative charge. (# valence e− in C) + (# valence e− in N) + 1e− for − charge  4 e− + 5 e− + 1 e− = 10 e− Exceptions to the Rule Octet Rule It is impossible to write good Lewis structure for molecules with odd numbers of electrons, yet some of these molecules exist in nature.

Less Than an Octet Aluminum and boron, which can function well with six valence electrons. Ex. BF3. Most elements to the left of the carbon group have so few valence electrons that they are in the same situation as boron: theyare electron deficient. In spite of these exceptions, Lewis theory remains a powerful and simple way to understand chemical bonding!. Resonance: Equivalent Lewis Structures for the Same Molecule • Consider: SO2 • Write the skeletal structure: OS O • Sum the valence electrons. (# valence e− in S) + 2(# valence e− in O) 6e− +2(6e− )=18e−

Valence Shell Electron Pair Repulsion (VSEPR) Theory Table 10.1 page 39-40 Predicting Geometry Using VSEPR Theory 1. Draw a correct Lewis structure for the molecule. 2. Determine the total number of electron groups around the central atom. 3. Determine the number of bonding groups and the number of lone pairs around the central atom. 4. Determine the electron geometry and molecular geometry. Representing Molecular Geometries on Paper Many chemists use this notation for bonds to indicate three-dimensional structures on two- dimensional paper.

The major molecular geometries used in this book are shown here using this notation:

Electronegativity • The ability of an element to attract electrons within a covalent bond is called electronegativity. • Oxygen is more electronegative than hydrogen, which means that, on average, the shared electrons are more likely to be found near the oxygen atom than near the hydrogen atom. Electronegativity Consider this representation of one of the two OH bonds:

The oxygen atom (getting the larger share) has a partial negative charge, symbolized by δ− (delta minus). The hydrogen atom (getting the smaller share) has a partial positive charge, symbolized by δ+ (delta plus). The result of this uneven electron sharing is a dipole moment, a separation of charge within the bond. Polar Covalent Bonds • Covalent bonds that have a dipole moment are called polar covalent bonds. • The magnitude of the dipole moment,and the polarity of the bond, depend onthe electronegativity difference betweenthe two elements in the bond and the length of the bond. • For a fixed bond length, the greater the electronegativity difference, the greater the dipole moment and the more polar the bond. Electronegativity • The value of electronegativity is assigned using a relative scale on which fluorine, the most electronegative element, has an electronegativity of 4.0.

• Linus Pauling introduced the electronegativity scale used here. He arbitrarily set the electronegativity of fluorine at 4.0 and computed all other values relative to fluorine. *page 48 * Identical Electronegativities • If two elements with identical electronegativities form a covalent bond, they share the electrons equally, and there is no dipole moment. • In Cl2, the two Cl atoms share the electrons evenly. This is a pure covalent bond. The bond has no dipole moment; the molecule is nonpolar. Large Electronegativity Difference • If there is a large electronegativity difference between the two elements in a bond, such as what normally occurs between a metal and a nonmetal, the electron is completely transferred and the bond is ionic. • In NaCl, Na completely transfers an electron to Cl. This is an ionic bond. Intermediate Electronegativity Difference • If there is an intermediate electronegativity difference between the two elements, such as between two different nonmetals, then the bond is polar covalent. • In HF, the electrons are shared, but the shared electrons are more likely to be found on F than on H. The bond is polar covalent. The Effect of Electronegativity Difference on Bon Type

The degree of bond polarity is a continuous function. The guidelines given here are approximate

Polar Bonds and Polar Molecules Does the presence of one or more polar bonds in a molecule always result in a polar molecule? The answer is no.

A polar molecule is one with polar bonds that add together—they do not cancel each other—to form a net dipole moment. If a diatomic molecule contains a polar bond, then the molecule is polar. For molecules with more than two atoms, it is more difficult to tell polar molecules from nonpolar ones because two or more polar bonds may cancel one another. Polar Bonds and Polar Molecules Consider carbon dioxide: Each bond is polar because the difference in electronegativity between oxygen and carbon is 1.0. CO2has a linear geometry, the dipole moment of one bond completely cancels the dipole moment of the other, and the molecule is nonpolar. Polar Bonds and Polar Molecules Consider water (H2O): • Each bond is polar because the difference in electronegativity between oxygen and hydrogen is 1.4. • Water has two dipole moments that do not cancel, and the molecule is polar. CHAPTER 12 The Gaseous State of Matter A gas is defined as a state of matter consisting of particles that have neither a defined volume nor defined shape. It is one of the fundamental states of matter A gas may consist of atoms of one element (e.g., H2, Ar) or of compounds (e.g., HCl, CO2) or mixtures (e.g., air, natural gas). Almost all elements or compounds can be a gas if it is hot enough. Example: Water: ice(solid) →water(liquid)→steam(gas) What is a Gas? Examples of gases at room temperature (25 0C and 1 atm pressure): Air (a mixture of gases), Ozone (O3), water vapor or steam, etc. 

There are 11 elemental gases (12 if you count ozone).

H2 - hydrogen N2 - nitrogen O2 - oxygen (plus O3 is ozone) F2 - fluorine Cl2 - chlorine He – heliu Ne - neon Ar – argon Kr - krypton Xe – xenon Rn - radon Properties of Gases 1. 2. 3. 4.

Have indefinite volume. Expand to fill a container. Have infinite shape – it will assume the shape of a container. Have low densities Ex. dair = 1.2 g/L at 25 °C dwater = 1.0 g/mL at 25 °C (Water is about 1000 times denser than water steam.) 5. Gas molecules have high velocities and kinetic energies To Define Gas: 1. Pressure (P) 2. Temperature (T) 3. Volume (V) 4. # of moles (n) #of gas particles PRESSURE Pressure: Force per unit area Pressure = force area Pressure depends on: 1) The number of gas molecules 2) Gas temperature 3) Volume occupied by the gas SI unit of pressure is the pascal (Pa) = 1 newton/meter2 Unit Conversions: 1 atm = 760 mm Hg = 760 torr= 101.3 kPa = 1.013 bar = 14.69 psi Pressure conversions Convert 740 mm Hg to a) atm and b) kPa

FACTORS THAT AFFECT PRESSURE 1. Number of gas molecules present (n) 2. Temperature (T) 3. Volume (v)

Pressure and number of gas molecules Number of Gas Molecules: - Pressure (P) is directly proportional to the number of gas molecules present (n)(at constant temperature (T) and volume (V )) Increasing n creates more frequent collisions with the container walls, increasing the pressure. P is directly proportional to n P1/n1=P2/n2 Pressure and Temperature - Pressure is directly proportional to temperature (at constant moles (n) and volume (V)) P=kT P1/T1=P2/T2  Increasing temperature causes: a) More frequent and b) Higher energy collisions Pressure And Temperature Gay-Lussac's Law: The pressure of a gas is directly proportional to temperature (at constant volume (V) and constant number of gas molecules (n). P1/T1=P2/T2 Pressure And Volume - Pressure is inversely proportional to volume (at constant moles (n) and temperature (T). Boyle’s Law Volume of the Gas Molecules: Pressure is inversely proportional to temperature (at constant moles (n) and temperature (T). P1V1=P2V2 What volume will 3.5 L of a gas occupy if the pressure is changed from 730. mm Hg to 600. mm Hg? SENSE CHECK: AS VOLUME DECREASES, PRESSURE SHOULD INCREASE.

Charles Law a) The volume of a fixed quantity of gas will increase if temperature of the gas is increased. b) The volume of a fixed quantity of gas will decrease if temperature of the gas is decreased.

Jacques Charles’ Law: The volume of a fixed quantity of gas is directly proportional to the absolute temperature of the gas at constant pressure. V1/T1=V2/T2 V1T2=V2T1 Charles’ Law Kelvin Temperature Scale As a gas is cooled by 10 C increments,the gas volume decreases in increments of 1/273. All gases are expected to have zero volume if cooled to −2730 C This temperature (−273 oC) is referred to as absolute zero. Absolute zero is the temperature (0 K) when the volume of an ideal gas becomes zero. All gas law problems use the Kelvin temperature scale!!! Formula: K=C0 +273 3.0 L of H2 gas at −15 oC is allowed to warm to 27 oC at constant pressure. What is the gas volume at 27 oC? Avogadro’s Law: Equal volumes (V) of different gases at constant T and P contain the same number of molecules/moles (n). Avogadro’s Law (another interpretation): At constant T and P, the number of gas molecules is directly proportional to the volume of the gas(es). n1/V1=n2/V2 Given the following gas phase reaction: N2 + 3 H2 2 NH3 If 12.0 L of H2 gas are present, what volume of N2 gas is required for complete reaction? T and P are held constant. By Avogadro’s Law, we can use the reaction stoichiometry to predict the N2 gas needed.

Given the following gas phase reaction: 2 H2 + O2

2 H2O

At constant T and P, how many liters of O2 are required to make 45.6 L of H2O? Sense Check: Less moles of O2 equal less L of O2!

Mole/Mass/Volume relationships Molar Volume: the volume 1 mol of gas occupies at STP molar volume = 22.4 L/mol at STP For chemistry, IUPAC established standard temperature and pressure (informally abbreviated as STP) as:

A temperature of 273.15 K (0 °C, 32 °F) and Absolute pressure of 101.325 kPa (14.7 psi, 1.00 atm, 1.01325 bar). The volume of O2 at STP is 44.8 L. What is the number of moles of O2? What is the mass (in grams) of 44.8 L of O2?

Combined Gas Laws A combination of Boyle’s and Charles’ Laws. Used in problems involving changes in P, T, and V with a constant amount of gas.

The volume of a fixed quantity of gas depends on the temperature and pressure.It is not possible to state the volume of gas without stating the temperature and pressure. Standard Temperature and Pressure (STP): 0.00 °C (273.15 K) and 1 atm (760 torr) What is the volume at STP for a gas that occupies 1.62 L at 616 torr and 42 °C?

Ideal Gas Law A single equation relating all properties of a gas.

PV=nRT Where R is the universal gas constant R is derived from conditions at STP. Calculate R. Knowns: P = 1.00 atm V = 22.4 L T = 273 K Solving for R R=PV/nT Calculate:

n = 1.00 mol

Units are critical in ideal gas problems! Ideal Gas Law – Molar Mass The ideal gas law can also be written in terms of molar mass of a gas. PV = nRT Kinetic Molecular Theory A general theory developed to explain the behavior and theory of gases, based on the motion of particles. Assumptions of Kinetic Molecular Theory (KMT): 1) Gases consist of tiny particles. 2) The distance between particles is large when compared to particle size. The volume

occupied by a gas ismostly empty space. 3) Gas particles have no attraction for one another. 4) Gas particles move linearly in all directions, frequently colliding with the container walls or other particles. 5) Collisions are perfectly elastic.No energy is lost during collisions. 6) The average kinetic energy for particles is the same for all gases (regardless of molar mass) at the same temperature. KE = 1/2mv 2 Where m is the mass and v is the velocity of the particle The average kinetic energy is directly proportional to temperature (in K). Gases which behave under these assumptions are known as ideal gases. Real Gases Real gases typically behave like ideal gases overa fairly wide range of temperatures and pressures. Conditions where real gases deviate from ideal gases: a) At high pressure (small volumes) Distance between particles is small and the particles do not behave independently. b) At low temperature Particles experience intermolecular interactions. Gas Laws: Relating pressure, moles, temperature, and volume of gases. 1. Boyle’s Law2. Charles’ Law3. Avogadro’s Law4. Gay-Lussac’s Law5. Combined Gas Law6. Ideal Gas Law7. Real Gas Law (not covered here) Dalton’s Law of Partial Pressures The total pressure of a mixture is the sum of the partial pressures of the different gases in the mixture. Ptotal =P1 +P2 +P3... Each gas behaves independently in the mixture. Application of Dalton’s Law Gases collected over H2O contain both the gas and H2O vapor. Vapor pressure of H2O is constant at a given T. Pbottle is equalized so that Pbottle = Patm thus Patm = Pgas + PH2O A 250. mL sample of O2 was collected over water at 23 0 C and 760 torr. What volume will the O2 occupy at 23 oC when PO2 is 760. torr? The vapor pressure of water at 23 oC is 21.2 torr. Gas Density Density of a liquid or solid is expressed in g/mL, but gas density is very low, so the standard units are g/L. The density of a gas at STP can also be related to the compound’s molar mass.

Note: gas densities must be cited at a specific temperature as volume changes as a function of temperature (Charles’ Law). Calculate the density of Cl2 at STP. Sense check: GAS DENSITIES are expected to be low.

Gas Stoichiometry

At STP: the molar volume can be used as a conversion factor to convert between moles and volume. Non STP Conditions: use the ideal gas law to convert between moles and volume. Given the following reaction: 2NaN3 (s)

2Na(s)+3N2(g)

If an air bag should be filled with a pressure of 1.09 atm at 22 oC, what amount of solid NaN3 is needed to filla bag with a volume of 45.5 L?...