Preparatory Notes

CHEM 142

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Table of contents

1. Conversions
2. Matter
   1. Organization of the Periodic Table
   2. Standard Chemical and Physical States of Elements
   3. Law of Constant Composition
   4. Law of Multiple Proportions
   5. Common Laboratory Separation Techniques
   6. Substance Classification
   7. Distinguishing Chemical and Physical Change
3. Atoms, Ions, and Molecules
   1. Parts of an Atom
   2. Molecular vs Empirical FOrmulas
   3. Naming Binary Ionic Compounds
4. Thermochemistry
   1. Calculating Kinetic Energy
   2. Calculating Specific Heat Capacity
   3. Heat Capacity Equation
5. Stochiometry
   1. Writing a Chemical Equation from a Reaction Description

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Conversions

\[1 \text{ kcal} = 4.184 \text{ kJ}\] \[1 \text{ cm}^3 = 1 \text{ milliliter}\] \[\left(a\cdot \frac{9}{5}\right)+31^\circ \text{F} = a^\circ \text{C} = a+273.15 \text{K}\]

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Matter

Organization of the Periodic Table

General Periodic Table:

Standard Chemical and Physical States of Elements

The standard state of an element is the chemical and physical state of the element at a pressure of 105 Pa (a pressure very close to 1 atmosphere) and a standard temperature, usually 25°C.

In other words, the metals except for mercury (Hg) are solids in their standard state. Mercury is a liquid. Most nonmetals are written as simple solids, except for the noble gases, the seven diatomic nonmetals, phosphorus, and sulfur. The diatomic nonmetals are mostly gases, but iodine is a solid and bromine is a liquid.

(The following elements are diatomic in their standard states: H, N, O, F, Cl, Br, and I.)

Law of Constant Composition

In a pure substance, chemical elements always appear in the same proportions.

Law of Multiple Proportions

The Law of Multiple Proportions is an empirical law about the composition of chemical compounds, proposed by the English chemist John Dalton (1766-1844) in about 1806. It says that if two elements can combine to form more than one compound, the masses of one element that combine with a fixed mass of the other elements must be whole number multiples of each other.

Common Laboratory Separation Techniques

• Filtration: passing a mixture through a filter, which traps solid components and lets liquid or gaseous components pass through. Filtration is most useful for separating heterogeneous mixtures of solids and liquids, or solids and gases.
• Distillation: heating a mixture until it boils, then collecting and cooling the vapor until it condenses. The condensed liquid is called the distillate. The distillate will have a higher percentage of the components with lower boiling points. Distillation is most useful for separating homogeneous mixtures of liquids that have widely separated boiling points.
• Extraction: putting a mixture in contact with a solvent in which only some of the components are soluble. The soluble components will be drawn out of the mixture by dissolving in the solvent. Extraction is most useful for separating solid or liquid mixtures in which some of the substances have significantly different physical properties so that they are soluble in very different types of solvent.
• Chromatography: flowing a mixture of liquid or gaseous components through a porous material called an adsorbant. Because the molecules from which the substances are made will wiggle through the tiny pores in the adsorbent at different rates, the substances in a mixture will spread out as they travel through the adsorbant. Chromatography works well with even very small sample sizes, so it’s a favored method for detecting trace (“very small”) amounts of substances.

Substance Classification

Hierarchy of substance classification:

• Substances
  • Pure Substance
    • Element (iron, nitrogen)
    • Compound (water, table salt)
  • Mixture
    • Homogenous Mixture (seawater, coffee, air)
    • Heterogeneous Mixture (milk, smoke, butter)

Determining compound classification:

• Heterogeneous mixture if a substance is made up of chunks, strands, layers, or blobs of different substances, homogeneous mixture otherwise.
• Element if a substance is made of just one type of atom, substance if there are 2+ atoms but atoms appear in a particular combination, heterogeneous mixture if atoms or molecules appear in any combination.
• Liquid homogenous mixtures are solutions. To determine if a substance is a heterogeneous mixture

Distinguishing Chemical and Physical Change

• Principle of Conservation of Mass: matter can’t be created or destroyed by ordinary events. (“Ordinary” events are events that don’t involve nuclear reactions.)
• Chemical Change: A change that re-arranges atoms into different molecules.
• Physical Change: A change that only moves atoms or molecules around, and doesn’t change the number or type of molecules.

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Atoms, Ions, and Molecules

Parts of an Atom

Molecular vs Empirical FOrmulas

• The molecular chemical formula lists every atom that’s in each molecule.
• The empirical chemical formula lists the ratios of the atoms in each molecule.
• The molecular formula \(C_4 H_4 O_2\) can be rewritten as the empirical formula \(C_2 H_2 O\).

Naming Binary Ionic Compounds

All binary ionic compound names follow this pattern: [cation name] [anion name].

Stock notation: You should only write the cation charge in a cation name when the metal can form more than one cation. Most of the elements that can do so are transition metals. You should leave the Stock notation out of any cation formed from an atom of a Group 1A or Group 2A metal, or from atoms of the elements aluminum (Al), zinc (Zn), and cadmium (Cd).

If an anion is an atomic ion of a nonmetal element, then its name is formed from the name of the nonmetal element by replacing everything after the first syllable with -ide.

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Thermochemistry

Calculating Kinetic Energy

Kinetic energy is energy in the form of motion. Everything that moves has kinetic energy, and anything that has kinetic energy must be moving. Qualitatively you should remember that kinetic energy increases with both mass and speed. In other words, the faster an object is moving the greater its kinetic energy will be. Similarly, the greater an object’s mass is, the greater its kinetic energy will also be.

\[E_k = \frac{1}{2}mv^2\] \[E_k := \text{ kinetic energy in Joules, or } \text{kg} \cdot \left(\frac{\text{m}}{\text{s}}\right)^2\] \[m := \text{ mass in }kg\] \[v := \text{ speed in }\frac{\text{m}}{\text{s}}\]

Calculating Specific Heat Capacity

The specific heat capacity of a substance is the amount of heat needed to raise the temperature of a gram of a substance by a kelvin. You can think of specific heat capacity as something like the “price” of a kelvin of temperature rise. It takes more heat to “buy” a kelvin when a substance demands a high “price” (has a high heat capacity) than when it demands a low “price” (has a low heat capacity).

\[c \text{ (heat capacity)} = \frac{Q}{m \cdot \Delta T}\] \[Q := \text{total heat used}\] \[m := \text{mass of pure sample}\] \[\Delta T := \text{change in temperature}\]

Heat Capacity Equation

\[\Delta E = m\cdot c \cdot \Delta T\] \[\Delta E := \text{thermal energy a substance absorbs or loses}\] \[m := \text{mass}\] \[c := \text{heat capacity}\] \[\Delta T := \text{change in temperature}\]

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Stochiometry

Writing a Chemical Equation from a Reaction Description

1. Identify all compounds and elements that take part in the reaction.
2. Decide which compounds and elements are reactants and which are products.
3. Write an unbalanced chemical equation.
4. Balance the equation.

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