CHM110 - Chemical Principles 1

Nanomaterials, Nanotechnology, Nanoparticles Macroscopic - everything we can see (millions of molecules together) Microscopic - single molecules Nanoparticle - contain 1000 atoms or molecules and can be seen under TEM Property changes when becomes nanoparticle.

Unit One: Lecture Notes Atoms # (protons) #e (electrons ) # atomic number ## mass number For equal # the # identifies the isotope Atomic Mass Mass Number Mass of atom (amu) Mass of atoms (g) amu/g Mole particles Shorthand Notation for the Atom Components of the Atom Percent Abundance (where is the % abundance of ) Molecules Formed when e- of one atom simultaneously is attracted to its own nucleus and another surrounding nucleus. Example When ➀ and ➁ move closet, attracted

Units Significant Figures - Refers to digits that were measured when rounding calculated numbers we pay attention to significant figures, so that we do not overstate the accuracy of our answers. Example - Better measurement, more accurate - Less accurate Accuracy - Proximity of a measurement to the true value of quantity. Precision refers to proximity of several measurements to each others. What Exists? Atoms Atomic Number - Number of protons in the atom Mass Number - Protons a

Micro - Only be seen with a microscope Macro - Can be seen with bare eyes Volume and - most commonly used Uncertainty in Measurement Use burrette for accurate (to the decimal) measurements Use graduated cylinder for precise (rounded) measurements Accuracy vs Precision What Exists? Atoms number of protons = number of electrons = atomic number number of protons + number of neutrons = mass number mass of 1 atom isotopes Atoms of the same element with different number of neutrons =

Boyle’s Law Pressure of a fixed mass of gas is inversely proportional to its volume at constant temperature. Initial Conditions Final Conditions at constant and Why? Constant value because moles and temperature are constant Proportionate to volume If volume increases, then as is constant Charle's Law The volume of fixed mass of gas is directly proportion to it’s Kelvin temperature at constant pressure Initial Conditions Final Conditions at constant and

Ex: Chlorine has two isotopes and . is 75%. What is the atomic mass of ? Ex: Uranium has atomic mass of 238.03 and it consists of 2 isotopes with mass of with mass of What is the % abundance of in naturally occurring uranium? What Exists? Matter Pure - Constant composition Mixtures - Variable composition Homogeneous (single phase) Heterogeneous (more than once phase)

Nucleus Binding energy released when nucleus is formed Nucleic energy A large amount of energy is produced by binding energy 1. Example Chlorine has 2 isotopes: 3s Cl and 37 Cl. Cl is 75%. What is the atomic mass of Cl. 35x75 = 262.5 37x25=935 => 925+262.5 =35.5 100 2. Example Uranium has atomic mass of: 238.03 It consists of 2 Isotopes 235: U with mass of 235.044, and 238U with mass of 238.05^ X: % of 238 U 238.051* X + (100-x) 235.044 = 238.03 * 100 238.051X + 2350YYX

Laws of Thermodynamics 1. Where: hear absorbed by the system work done on the system 2. for all spontaneous processes 3. entropy of a perfect crystal at Enthalpy change for a process Heat absorbed by the system when the process is carried out at constant pressure final stateinitial state If Exothermic if Endothermic Calculating/Measuring Carry out the reaction in an isolated system Measure T due to reaction Calculate H from & What is ?

Separation Techniques of mixtures Qualitative analysis Separation and identification of the components of the mixture, what is present. Quantitative analysis Determination of the quantity of each component of the mixture. how much is present. Crystallization - Separation of solids from a liquid solution. A crystal is an ordered arrangement of molecules. First formation of a crystal needs another material for the crystallization to happen (even dirt). Evaporation of water for example to form a N

Qualitative - Description with words Quantitative - Quantity shown with numbers Separation Techniques Crystallization Distillation Chromatography more affinity to the solvent travel more/faster/eluted first highest on the paper Solvent extraction Calculating Moles Mole is a number # of mols = mass in g Ex: How many atoms are in of copper? Ex: How many grams of will precipitate out if reacted? Combustion Reaction Combustion - Substance reacts with to form and Ex: Upon comb

Compressibility (P, V, T) Boyle's Law: (constant T and n) Charles Law: (constant T and n) Gay-Lussac: Ideal Gas Law more collision between molecules at different pressure relationship between and change, gas constant Kinetic Molecular Theory Molecules are far apart; most of the volume is empty. Molecules are in constant random motion they travel in straight line paths. Molecules exert no force on each other except during collisions. Average kinetic e

Find the Empirical Formula of Vitamin C Example A 6.49 mg sample of ascorbic acid which is known to contain only and was burned in a combustion apparatus. The increased weights of the absorption tubes showed that 9.74 mg of and 2.64 mg of were formed. What is the empirical formula of the acid? 6.49 mg 9.74 mg 2.64 mg Moles of Mass of compound = number of moles atomic mass # moles of Find the % Composition of a Mixture Example A mixture of and with a mass of 4.000

Moles Units mole of oxygen atoms = atoms = mass of grams mole of carbon atoms = atoms = mass of grams mass of carbon atom = Calculation of Moles How many moles in 36g (C) Carbon? # How many moles of NaOH are there in 8g? = Calculate the # of atoms in mol of Copper. How many atoms are there in g of Carbon? A silver ring contains silver atoms. How many moles of silver are in a ring? # How many grams of Ag Cl will precipitate out if mole is reac

Solubility Rules What Happens? Reactants and products? number of reactants number of products Elements and compound Neutralization/acid-base reactions Acid - Produces from Base - Produces from Solubility Guidelines soluble Halides insoluble Most salts of and Important Redox Reaction Oxidation - Loss of electrons Reduction - Gain of electrons Oxidizing Agent - Causes an element/compound to lose electrons and gets reduces Reducing Agent - Causes an element/compound to

Principles of Chemical Equilibria Definition Rate - Change in [ ] / time at equilibrium K - equilibrium constant k - rate constant - # moles of gaseous product - # moles gaseous reactant Factors Affecting 1. temperature Exothermic : , (shift left) , (shift right) b. Endothermic : , (shift right) , (shift left) Van Hoff Equation assume derived from it, NOT the equation Quantitative Exothermic (negative #) as

Combustion - A reaction in which a substance reacts with oxygen to form oxygen containing compound and heat. Example - Combustion of natural gas (methane) produces carbon dioxide and water Empirical Formula = Molecular Formula Ratio of atoms in a molecule - Actual number of atoms molar mass/empirical formula molar mass Example 1 - Aspirin Given - % % % Find - Empirical formula Moles of each element - Mass of Mass of Mass o

DEFINE a) ACIDS AND BASES 1. arrhenius (has to be in water) Acid - proton donor, will donate , yields Base - yield when dissolved in 2. bronsted-lowry (does not have to be in water) Acid - proton donor, donate Base - proton acceptor 3. lewis Acid - accepts pair of Boron will accept pair of Base - donates pair of can donate that pair to make a bond 4. solvo-system Acid - Generates the cation form of solvent in solution Base - Generates the anionic form of solvent i

Balancing Reactions Balancing Just add Example removed since remained throughout the reaction Example

Test for Present Chloride Using Silver nitrate, if there is white precipitation, there is Chloride. precipitation reaction mass % of each component of original mixture mass mass of mass mass of Solve for Mass of divalent = Calcium (by the valency) Metal Chloride = % of Chemical Reactions 1. Classify by number of reactants and products A. Number

Concentrations of Solutions % w/w - g solute/100 g solution % w/v - g solute/ 100 ml solution % v/v - ml solute/100 ml solution Example 20 ml of a 5% solution of can completely neutralize 20 ml of a 5% solution of . True or False? False → Since we don't know if the number of mole is equal to each other or we don't know how many moles are involved. molar solution Concentration expressed in mol Molarity (M) = moles of solute / L solution Example What is the molarity of a 40 moles of dissolve

Buffer Solutions A solution containing a weak acid and it's anion in approximately equal amounts Or weak base and its cation The pH of the solution will be close to the pKa of the acid component of the solution Weak acid or cation of the weak base A buffer solution has the ability to maintain its pH approximately constant, despite the addition of acid or base to it. Add 0.01 mole to of... pH changes pH changes pH changes pH changes pH doesn't change much buffer

Bleach - Sodium Hypochlorite in water Danger Bleach and Ammonia Toxic Example Oxidation - Loss of electrons Reduction - Gain of electrons took the electron oxidizing agent caused something to gain electrons reducing agent. Spectator Ion - Does nothing stays the same, doesn't go threw oxidation or reduction Metals Reducing Agent. Strongest Weakest Redox Reactions One part of the molecule is positive and the other negative each 0 i

Chapter 9, 10.1, 10.5-10.11, 13.8, 11.1-11.7 9.1 THE NATURE OF ENERGY Energy - capacity to do work or to produce heat Law of Conservation of Energy - energy can be converted from one form to another but can be neither created nor destroyed Potential Energy - energy due to position or composition. Attractive and repulsive forces also lead to it Kinetic Energy - due to the motion of the object and depends on the mass of the object (m) and its velocity Energy can be converted from one form to an

Compressibility (P, V, T) Boyle's Law Constant T and n Increase of pressure = decrease of volume Charles Law Constant P and n Gay-Lussac dry ice = solid Temperature is constant Absolute Zero = lowest temperature you can obtain at How close can the volume go to zero every gas when cooled down closely, it becomes a liquid Liquid Helium - Closest gas that can go to almost 0 Why is absolute zero ? - is the volume of a gas at - is volume of at any temperature Volume of any ga

+3 -3 Add 2+3 2. 3. Add 2 + 3 Chemical Reactions Soluble/Insoluble Precipitation Reagents Phrases In Periodic Table At room temperature, all elements are solids except → Mercury and Bromine are liquids Hydrogen, Nitrogen, Oxygen, Florine and the Noble Gases are gases in r

Importance - These are two strong oxidizing reagents. are two strong oxidations Each Carbon is +3 Oxidation Reduction Potassium Dichromate Concentrations of Solutions Physical Method % w/w g solution/100 g solution Solvent %w/v g solution/100 mL solution Solute %v/v mL solution/100 mL solution Solution Solution and liquid are different Solvent State & Solution Soda

What is Pressure? forces of gas particles striking wall / area Pressure - Volume Relationship Pressure in the higher altitude All the gases in the air think they're alone and applies pressure partial pressure = total pressures applies and is the same for all 3 gases Example Dalton's Law of Partial Pressures Example A compound contains only N and H and is 87.4% N by mass. A gaseous sample of the compound has a density of at 710 torr and Find the molecular formula

1M (of any substance) in 1L → 1Molar solution Example → 1L of the solution not 1L of water Together the substance and solute = solution M (Molarity) = Moles/V (L) # of moles of solution Volume of solution in litres Dilution Density Molarity → (M) moles solute / L solution Molality → (m) moles solute / kg solvent Normality → (N) equivalents solution / L solution Mole Fraction → moles solute / moles of solute + moles of solvent Example #1 Example #2 X= 0.05 p= 0.997 g/m Example #

question #3 question #4 35.0 mL Acetone 65.0 mL Water Volume = 97.0 mL The density of pure Acetone is 0.792 g/mL The density of pure Water is 1.00 g/mL Find M, m, and X acetone for the solution Why is Molality Always Greater than Molarity? For the same solution, calculation mathematically When making a solution of 1 molal, you take 1000g of Water and 40g of solute. The volume goes up and by definition Molarity = Moles/Litre Molarity is less than Molality When is Molarity a

Kinetic Molecular Theory Molecules are far apart; most of the volume is empty Molecules are in constant random motion; they travel in straight in paths Molecules exert no force on each other excepting during collisions The average kinetic energy of the molecules depends upon the temperature Exam Question What are the assumptions of the kinetic molecular theory? Elastic Collision no loss in kinetic energy Kinetic energy of gas only decided by temperature Pressure inside vs pressure outside Les

Ideal Vs Non-Ideal Higher Mass - Diffusion rate is lower Lighter Mass - Diffuses faster Lighter Isotope - Diffuses faster Separation of from Mean Free Path - Average distance, the particle travels between collisions with other moving particles Non-Ideal Gas Behaviour Molecules occupy volume which is jot negligible as compared to the volume of the container There are intermolecular forces between pairs of molecules Van der Waal's Equation What is excluded volume? Volume actuall

Pressure Increases, Volume Decreases Boyles Law P Dry ice (solid ) and liquid soap Charles Law V T Gay-Lussac VT°C(1+T°C/273) Absolute Zero Gas Heated Expands In reality the volume decreases by 273>, at some stage it becomes a liquid. At around -273.15 the gas becomes liquid This is the absolute 0 is the liquid of a gas at volume at any temperatures when Volume is proportional to the number of moles Depends on temperatures pressure V n at constant T an

Pressure Volume Relationship There is an increase in gas sample size at higher attitude. KE - Kinetic Energy M - Mass V - Velocity In the Experiment Done in Laboratory there are Two Forces (opposing) Gas and liquid in the tube trying to come out (beaker). Atmospheric pressure pushing the liquid outside the tube into the beaker. They are Two Equal Forces Acting on Each Other → Level Stays Partial Pressure - Hypothetical pressure in which one gas would occupy the entire volume of the entire mi

Reaction Dynamics Equilibrium - Reversible reaction Decomposition and Calcium Carbonate in an open vessel Lime → Naturally found in the earth/within shells of aquatic animals Reaction Dynamic reactant products Forward reaction and backward reactions are present until forward and backward rates are equal Rate is equal, nit the quantity (concentration) [Reactant] , [product] forward reaction rate decreases Eventually, the products can react to re-form some of the reactants Reverse reacti

Why does a Balloon Shrink in Liquid Nitrogen? Because when temperaturevolume and opposite. Maxwell Boltzmann Distribution Law Most molecules have the velocity somewhere in the middle, some have above and some below. How nature works? As the molar mass of the gas increases the velocity decreases. If M is high V is low Probability number of molecules having the same particular velocity. Highest Probability for M and Lowest for Oxygen Because molar mass is the lowest for Hydrogen. 1 mol o

Equilibrium Constants Concentration of a gas in a mixture Proportional to its partial pressure Equilibrium constant Ratio of the partial pressures of the gases OR Deriving the relationship between and → Moles of → Volume of gas substituting Used for For Heterogeneous Equilibria Pure Solids - Concentration does not change during the reaction Solids and Liquids - Not included in the equilibrium constant expression Example → Calculating Equilibrium Constan

Electrolytes What are Electrolytes? Conduct electricity either in the molten state or in solution Example Strong and weak electrolytes High degree of dissociation - large number of ions conduct electricity better Example Low degree of dissociation - few ions = poor conductivity Example Strength of an Electrolyte - Inherent property not based on the concentration Sugar/glucose electrolytes Arrhenius Concepts Acids and bases = electrolytes Acids dissociate in give ions Bases dis

The Kinetic Theory - Mean free path of a particle such as a molecule is the average distance the particle travels between collision with other moving particles. Real Gases Obey Actual Volume Ideal Pressure Van der Waals Equation - Force of a Att

When gases behave non-ideal High pressure and low volume, intermolecular distance can become shorter When gases behave ideal Low pressure and high volume High temperature van der waal's equation The value of the Van Der Waals constant is a measure of the deviation from ideal behaviour Boyle's temperature For every gas there is a temperature, the Boyle temperature at which that gas behaves ideally For one mole of a gas or We find that at high temperatures the behaviour is mor

Acid Ionization The ability of an acid to lose a proton Larger the value of the acid ionization constant , higher that in the solution Stronger the acid When you know the ionization constants of acid, you can get the relative strength of different acids at a particular temperature Example Small the reaction will shift to right with very limited excess Example → Is a weak acid, inorganic acid When comparing and , is a stronger acid due to the ionization constant is larger tha

(1) Kp 1 = 0.212 atm (2) (Reverse reaction and then divided by 2) Knew = What happened to the equation? apply reverse reaction rules. at Find for at If you preform a reaction in an aqueous Solution you put in the molarity of the solution (In terms of molarity) Kc If you preform a reaction where all the reactants and products are gases, then you reprise the concentration with another quantity: Kp (partial pressure) proportional to concentration.

Identify the Acids Arrhenius Bronsted-Lowry Lewis Acid can behave as a base depending on the solvent Example most acidic Scale Example If an acid has an concentration of , find the Increasing ions are decreasing Decreasing ions are increasing What is the of ? of Water Autoionization of Reaction is endothermic so at low temperature becomes lower as dissociation is decreased Temperature is lower, is high Temperature is hig

- has no effect Equilibrium Constant - When temp. is constant, independent of the Initial amounts of reactants and products. Constant of Product / reactant. Equilibrium Constant - K: will not depend on the initial concentration of the reactant and product. It will depend on temperature only. Equilibrium Constant: Note: Initial Concentration Change in concentration Equilibrium concentration Reaction Quotient - Concentration ratio of the product (raised to the power of thei

Acid and Base Calculations → [acid] → [base] → Volume of acid → volume of base of is added to of . What is of the solution? Species = Mass balance = Charge balance = From step (2)

The Effects of Temperature on Equilibrium Position: Exothermic reactions - release energy and endothermic reactions absorb energy. Heat Produced = exothermic reaction Heat reactant = endothermic reaction Le Chatelier Principle to Predict the Effect of Temperature Changes Increasing the temp. is like adding heat. Le chatelier Principle: equilibrium will shift away from added heat. Adding heat to an exothermic reaction will decrease the concentrations of products and increase concentrations of re

PH Indicator shows a colour when is equal to or less than and show a different colour when is equal to or higher than Colour change occurs when changes from to Buffer Mixture of a weak acid and conjugate salt approximately in the equal amount Keeps the same even if a small amount of acid/base was added Acid-Base Indicator Indicators mark the end point Used to detect the equivalent point of a titration, must only add a very small amount of indicator Indicators are weak acids/bases a

Binary Acids: Molecular compounds in which hydrogen is combined with a second nonmetallic element The strength of the binary acids increases as the electronegativity of increases. HCI is a stronger acid than which is a stronger acid than The greater the radius of the weaker the bonds and the acid is stronger. ( is stronger than stronger than ). The decrease in bond strength outweighs the decrease in electronegativity of , so the acids becomes better proton donor. Oxoacid: An