General Chemistry Study Guide PDF

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General Chemistry Study Guide Chapter 1 composition Structure Dynamics Energetics Icon Standards are different College is harder than high school Chemistry Scientific method - Begins with observation Communicating data - What are the results of my experiment - Does it make physical sense - Can i communicate them others - Are the results consistent? What is the error margin - How reliable is my measurement? What is a measurement? - Quantitative observation - Every measurement has a unit, number, an uncertainty SI Units What is energy - Energy is the capacity to do work - 2Na + Cl2 -> 2NaCl + heat Forms of energy - Kinetic – motion - Potential – position but not (yet) actualized - Chemical – chemical change - Thermal – heat - Electrical – flow of charge Units of energy - Joule (j) - si derived unite for energy - 1 J = 1 kg x m^2/s^2 - 1 calorie (cal) – energy to raise the temperature of 1 g of water by 1°C - Nutritional calorie (Cal) = 1 kcal - Kilowatt hr (kWh) – see electrical bill (1 kWh = 3.60MJ)

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1 kwatt × 3600 s = 3600 kJ

Exothermic & endothermic processes - During the reaction, heat can be released (a) or absorbed (b) as the reactants go to products - Exothermic – heat released during reaction - Endothermic – heat absorbed during reaction Heat and temperature - Temperature is the measure of this constant random motion (thermal energy). Heat is the measure of the thermal energy transfer (exchange of thermal energy). Heat capacity - amount of thermal energy (J) required to change a specified mass of given material by 1 oC - Specific Heat Capacity or Specific Heat - Heat capacity in units J / g oC, represented as C

Chapter 3 Matter - If it occupies space adnd ahs weight it is matter - All matter composed of atoms When atoms share connections, a chemical bond is formed When 2 or more atoms are bonded, a molecule is formed Molecular geometries contribute to properties bonds between atoms within a molecule are intramolecular (stronger) The bonds between molecules are intermolecular (weaker) Solids and liquids have fairly strong intermolecular forces that cannot be ignored. Solids and liquids are called CONDENSED phases. In gases, the intermolecular forces are virtually nonexistent and can often be ignored. States of Matter ● Solid particles are adjacent and locations fixed. Molecules may vibrate about location but do not move around each other. Fixed volume and shape. ● Crystalline – long-range, repeated order ● Amorphous – local but not long-range order ● Liquid (fluid*) particles are adjacent but move around each other. Fixed volume, shape morphs to container. ● Gas (fluid*) particles are separated and move around each other. Compressible (neither volume nor shape is fixed) Composition of matter Pure substance - composed of 1 type of atom ore molecule Element - Matter of one type of atom has elemental composition Compound - Compound material formed of two or more elements in definite proportions. Can be decomposed to elements and simpler compounds. Mixture - composed of 2 or more different types of atoms and/ or molecules

Mixtures Homogeneous - mixture composition uniform throughout the matrix Heterogeneous - mixture composition varies throughout the matrix Properties of matter Physical properties – observed without changes in chemical composition; □ includes phase transition (s↔l↔g) – odor (vapor pressure), color (absorbance), texture (state), melting & boiling points, density, reflectivity etc Chemical Properties - observed with changes in chemical composition; – Acidity, oxidation state, corrosiveness, toxicity, flammability etc. Chemical reaction ● Chemical properties require change in composition, a chemical change. ● A reaction is the process of chemical change. ● In chemical change, reaction forms a new and different material from the reactant. ● In a chemical reaction, matter is neither created nor destroyed. (law of conservation of mass) ● This forms the basis for chemical analysis and materials characterization.

Chapter 4 Matter atoms and elements First modern chemist - Robert boyle - 1627 - 1691 - The sceptical chymist (1661) - the birth of chemical analysis - Defined elements as pure substances that cannot be chemically decomposed into simpler substances The father of modern chemistry - Antoine lavoisier (1743-1794) - Elementary treatise on chemistry 1789 - the conservation of mass law - helped construct the metric system and develop chemical naming - the first list with 33 elements - burnt hydrogen in air to make water and decomposed it into H and O. John dalton (1766-1844) - A new system of chemical philosophy (1803) - law of multiple proportions - combined ideas of Boyle and Lavoisier - formalized Atomic Theory - published a table of relative atomic weights (hydrogen being 1) Dalton’s Atomic Theory 1. Atoms are the smallest piece of all elements. 2. Atoms are tiny indivisible and indestructible particles.

3. Atoms of a given element have the same mass & properties. Atoms of different elements have different masses and properties. 4. Atoms combine in simple whole-number ratios to form compounds. 5. Chemical reactions reassemble atoms to form different compounds Electrical charge ● A fundamental property of matter observed since ancient times. ● Think lightning, static electricity etc. ● “Electron” is Greek for amber. Discovery of nucleus ● An ‘alien’ student of J J Thompson ● Born in New Zealand ● Set out to experimentally prove the plum pudding model. ● Experimental evidence by Ernest Rutherford disproved the plum pudding model. ● Positively charged particles (α-particles) were shot at a thin gold foil. ○ Expected: a-particles (a speeding +2 charge) should pass right through the “pudding.” – ○ Observed: small fraction of α–particles were deflected Nuclear theory of an atom ● Positive charge must be concentrated in a small volume. ● 1. Most of atom’s mass & all positive charge is in a “nucleus” ● 2. Most of the volume of atom is empty space with the tiny negative electrons dispersed around the nucleus. ● 3. The atom is electrically neutral. The total positive charge in nucleus equals the total negative charge of the electrons. Nucleus ● Nucleus is very small (most of the atom is empty space) ● Positive charge is concentrated in the nucleus ● Nucleus contains 99.9% of atom’s mass density = 2 x 1014 g / cm3 • Electrons are compounds - Compound - material formed of two or more elements in definite proportions. Compounds can be decomposed to elements and/or simpler compounds. Mass of an element in a compound - Chemical formula gives the whole-number ratio of moles of element in the compound Mass percent of an element - To calculate percentage we divide the part (mass of element) by the whole (mass of compound) - Mass percent composition or mass percent of an element is the element’s percentage of the total mass of the compound. Empirical formula - Empirical formula is the simplest whole-number ratio - Laboratory analysis often provide masses of each element in a compound, not chemical formulas

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Knowing the ratios between masses of elements we can obtain empirical formula for a compound.

Empirical formula vs. molecular formula - Molecular Formula: C2H4O2 Acetic acid (vinegar) - Specific # atoms in a compound. - Empirical Formula: C1H2O1 Smallest whole-number ratios (many compounds)

Chapter 7 Chemical Reaction - Forms a new and different material from the reactant(s) - Atoms are rearranged - Easily identifiable at the molecular level Evidence of a Chemical Reaction - Change in color - Formation of gas - Light emission - Emission or absorption of heat - Formation of a solid in a clear solution (precipitation) The Chemical Equation - Symbolical summary of what happens at the molecular level - Macroscopic: moles (and grams) of chemicals are produced - Molecular: atoms are rearranged - Symbolic: chemical notation is used to record what happens Balancing a Chemical Equation 1. Write the unbalanced equation with all chemical formulas. 2. Start with the most complex formula a. First balance elements that occur in only one other species. b. Usually metals before non-metals. c. Continue until all elements are balanced. 3. Change coefficients to smallest integer ratio. Why are aqueous solutions important? - Cytosol is the aqueous solution of ions, proteins, and nutrients surrounding organelles within cells - Aqueous solution is a homogeneous mixture of a substance with water - Solution of sodium chloride in water is called a saline solution - 5% solution of acetic acid in water is called white vinegar - 3-6% solution of sodium hypochlorite (NaClO) in water is called household bleach - When an ionic compound is dissolved in water it usually dissociates into ions (cations and anions) Electrolytes - Substances that dissolve to form conductive solutions are called electrolytes. These solutions contain dissolved cations and anions

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Substances that dissolve to form non-conductive solutions are called nonelectrolytes. These solutions contain neutral molecules (for example, glucose solution, ethanol solution). Weak electrolytes partially dissociate into ions and therefore the solutions only conduct electricity poorly (for example, water).

Solubility Rules - When an ionic compound is soluble in water (for example, LiNO3) it means this compound forms a strong electrolyte solution (Li+ and NO3 ions in solution). - When an ionic compound is insoluble in water (for example, PbS) it means this compound does not form a solution (stays primarily as PbS). Gas Evolution Reactions - If an aqueous reaction forms gas as a product, it is called a gas evolution reaction. REDOX Reactions - Oxidation-reduction reactions are driven by the transfer of electrons. - The element that becomes more positive is oxidized. - Zinc atoms lose electrons and are therefore oxidized. - The element that become more negative is reduced. - Copper(II) ions gain electrons and are reduced. - Examples: - Rusting - Bleaching - Combustion - Batteries - Fuel cells - Citric acid cycle - Photosynthesis - Solar cells Combustion - A type of redox reaction of a substance (fuel) with O2 Fuel Cells - A hydrogen fuel cell can be used to power an electrical device using the reaction: - 2H2 + O2 → 2H2O + q - electrons are donated by hydrogen - 2 H2 → 4H+ + 4 e- and accepted by oxygen - O2 + 4 e- → 2 O2- How can hydrogen be obtained for this reaction? - By using a battery, electrons can be “pushed” to transfer in the opposite direction driving the reverse redox reaction: - 2H2O → 2H2 + O2 - Both H2 production and consumption are redox reactions. How To Recognize REDOX? - To identify oxidation-reduction reactions at this point:

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Look for change in ionic charge (electron transfer) between substances. Redox reactions are often between: - elemental metal and nonmetal(s): - Pb(s) + S(s) à PbS(s) (galena) - 4 Fe(s) + 3 O2(g)à 2 Fe2O3(s) - substance and elemental oxygen (combustion): - C3H8(l) + 5 O2(g) à 3 CO2(g) + 4 H2O(g) - elemental metal and salt or acid: - Zn(s) + CuSO4(aq) à ZnSO4(aq) + Cu(s) - 2Al(s) + 6HCl(aq) à 2AlCl3(aq) + 3H2(g)

Types of Reactions - By types of chemistry: - Precipitation (solid formation) - Acid-base (H+, OH-) - Gas evolution (bubbles!) - Redox (electrons exchanged, combustion) - The second classification scheme is by atom/group of atoms exchange pattern: - Double displacement - AB + CD -> AD + cb - Single displacement - AB + C -> CB + A - Synthesis (combination) - A + B -> AB - Decomposition - AB -> A + b - A, B, C, D are atoms of different elements or groups of atoms (such as polyatomic ions) Displacement Reactions - AB + C -> CB + A - Single displacement - Zn(s) + CuSO4(aq) à ZnSO4(aq) + Cu(s) - 2Al(s) + 6HCl(aq) à 2AlCl3(aq) + 3H2(g) - AB + CD -> AD + CB - Double displacement - acid-base, precipitation reaction and more Synthesis Reactions A + B -> AB Decomposition Reactions - AB -> A + B - The tabletop “volcano”: - Oftentimes decomposition reactions require additional energy (in the form of heat, light, or electrical energy) to proceed. - For example, decomposition of water (electrolysis of water) occurs when electrical current is applied: - 2H2O 2H2 + O2

Chapter 8 Stoichiometry - the study of the numerical relationships between chemical quantities of reactants and products in the chemical reaction. - Everything is consumed in the reaction as written down in the form of a balanced chemical equation Mass of a Compound Produced or Consumed in a Reaction - 2 conversion factors needed: - Molar mass is one conversion factor. - Molar ratio from a balanced equation is another. Theoretical vs Actual Yields - The amount of product calculated from the equation (taking into account limiting reagent) is the theoretical yield. - When you run a reaction, the amount of product collected is the actual yield. - These may not be the same because a fraction of the product may spill, escape as gas, stay deposited on the walls of the test tubes etc. - There is also such thing as human error ... - The percent yield is the amount of the actual yield compared to the theoretical yield. - actual/theoretical x 100% = percent yield Enthalpy of a Reaction - In a chemical reaction, energy is gained or lost. - Products may contain less or more energy than reactants. - Energy produced or consumed in the reaction is called the enthalpy of reaction: ΔHrxn. - Exothermic - heat released to surrounding - Endothermic - heat absorbed from surroundings Exothermic Process - A type of reusable instant HOT packs produces heat by crystallization of sodium acetate. - Precipitation is initiated by compressing a metal disk within the pouch of supersaturated NaCH3COO solution. The process is exothermic and the pouch warms up as the salt crystallizes. Enthalpy and Chemical Equation - Enthalpy of reaction (ΔHrxn) follows molar ratios of the balanced equation and is listed to the right of the equation. - Combustion of iron metal is an exothermic reaction that releases 1652 kJ of heat per mole of balanced reaction: - Energy is produced by the reaction, so the sign for the enthalpy is negative.

Chapter 9 What vs Why - The Periodic table allows us to predict the properties of an element based on its position in the table and empirical patterns. - It does not explain why the pattern exists. - Quantum Mechanics is a theory that explains why the periodic trends in the properties exist.

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To understand quantum mechanics a good place to start is with the study of light and waves.

Properties of a Wave - The amplitude is the height of the wave - the amplitude is a measure of intensity (brightness) - the larger the amplitude, the brighter the light - The wavelength, λ, is a measure of the distance, in meters, between the maxima of the wave - the distance from one crest to the next - The frequency, ν, is the number of waves that pass a fixed point in a given period of time - the number of waves = number of cycles - units are hertz, Hz, or cycles in one second = s-1 - 1 Hz = 1 s-1 - The speed of light, c, is the number of meters light travels in one second Inverse Relationship of Wavelength and Frequency - Longer wavelengths have lower frequencies - (towards the red side of the spectrum) - Shorter wavelengths have higher frequencies - (towards the blue side of the spectrum) - Different types of electromagnetic radiation have different wavelengths and frequencies but all have the same speed. Black Body Radiation - Classical physics predicted unlimited energy emitted by a heated black body. - Actually, even at extremely high temperatures very little of UV radiation is produced. - y. 1911 “Ultraviolet catastrophe” - Max Planck solved the problem with the assumption that electromagnetic energy is emitted/absorbed as discrete quanta of energy. - Nobel prize in 1918. Photoelectric Effect (y. 1921 Nobel Prize) - In 1905, Albert Einstein elaborated on the idea that light (electromagnetic radiation) is delivered in ‘packets’ called ‘quanta’ or photons. - Nobel prize in 1921. - The energy of a single photon is directly proportional to its frequency and inversely proportional to its wavelength. - The proportionality constant is called Planck’s Constant, h - h = 6.626 × 10-34 J·s for a single photon - c is the speed of light, 3 × 108 m/s y. 1913 Bohr’s Atomic Model 1. Electrons in an atom can only occupy certain orbits corresponding to certain discrete energies. 2. Electrons in permitted orbits have specific, “allowed” energies; these energies will not be radiated from the atom (“stationary states”). We say that the energy is quantized. 3. Energy is emitted or absorbed by the electrons only as the electrons change from one allowed energy state to another. The energy is emitted or absorbed as a photon, ΔE=hν.

Bohr’s Orbits and Energy - Electrons in an atom can have certain discrete energy levels depending on the orbit they occupy. - Each orbit is located at a fixed distance from the nucleus. - The energy of each orbit is characterized by an integer—the larger the integer, the more energy an electron in that orbit has and the farther it is from the nucleus. - The integer, n, is called a quantum number. - n=1 is called the ground state; - n = 1, 2, 3, ... and on are the excited states. - Bohr’s model successfully explained emission lines for atomic spectrum of hydrogen. Particle-Wave Duality - We have seen how light has characteristics of both a particle (photon) and a wave. The same is true for an electron. - Need to switch from thinking of fixed orbits (in Bohr’s model) to adopting the idea of electron-waves, or quantum-mechanical electron orbitals. - In 1926, Erwin Shrödinger derived a complex equation incorporating the dual nature of an electron: wave and particle. - Nobel prize in 1933. Schrodinger Equation - Quantum mechanics mathematically describes the ‘wave’ or ‘wave functions’ of electrons in an atom. - Statistical in nature. - The orbits are replaced with orbitals but labeling (n = 1, 2, 3, ...) and energy levels are the same. - Visualized as probability density maps Probability Density Map - We can calculate probability of finding an electron at each location. - If we model electron position over a very long time, we can obtain a probability map. - The shape of an electron orbital will then be the 90% probability cut off surface. - A particular electron is within the volume 90% of the time. Principal Quantum Number - In Schrödinger’s wave equation, there are 3 integers, called quantum numbers, that determine the wave or orbital. (the 4th, called spin, will come later) - The principal quantum number, n, specifies the main energy level for the orbital. (same as the Bohr model) - Positive integer (n = 1, 2, 3, 4, ...) - Commonly called principal shell - As the value of n increases: - The average distance of the e- from the nucleus increases - The energy increases (since the negative electron is more stable near the positive nucleus.) Other Quantum Numbers - In Schrödinger’s wave equation, An orbital is labeled by a set of three quantum numbers: 1. The principal quantum number, n, describes the energy level (of the hydrogen atom). (shell: 1, 2, 3, ...)

2. The angular momentum quantum number, l, describes the “shape” of the orbital. (sub-shell: s, p, d, f, g,...) 3. The magnetic quantum number, ml, describes “orientation” of an orbital in space. s-Orbitals - Spherical in shape - Radius of sphere increases with the value of n - In an undisturbed hydrogen atom, an electron is found at 1s orbital - (lowest energy state = ground state of atom) p-Orbitals - Dumbbell shaped - p-orbitals have 2 lobes and are often labeled with the axis along which they line up (x, y, or z) - X, Y, or Z direction is described by m, (magnetic quantum number) d-Orbitals - 4 of 5 d-orbitals have 4 lobes; the other resembles a p orbital with a donut around the center Energies of Shells and Subshells - As we fill shells with electrons, energy levels change. - For multielectron atoms, the subshells of a principal shell have slightly different energies. - s < p < d < f. - Each subshell contains one or moreorbitals: - s subshells have 1 orbital. - p subshells have 3 orbitals. - d subshells have 5 orbit...