Chem midterm 1

second ionization energy

(now have a cation so much harder to do) Always harder to remove a 2nd electron than a first bc now more protons than electrons It is hardest to remove a 2nd electron when it comes from a lower shell. - For example, it is much harder to remove a 2nd electron from Na than from K.

concentrations

Concentrations are measured in mol/ liter= M Concentration imp when talking about solutions Molarity is the number of particles per some unit volume This is called molarity The water is called the solvent The solute is the species dissolved in the water **you can determine the number of moles if you know the concentration and the volume— interconvert molarity volumes and number of moles Ex: 50 ml of a 2 M solution contains (2 mol/L)— (.05 L)= .1 mol of solute

polar v non polar (ELECTRONEG DIFF PUT ON SHEET??) also melting point/boiling point low or high conduct electricty as solid or aqueous? dissolves in polar solvent? dissolves in nonpolar solvent?

Electronegativity difference Polar .5-1.7 Non 0 - .4 Melting point/boiling point Polar Low Charges on molecule can interact w charges on other to make bond stronger when becomes polar Non Very low- bc don't have strong forces bw elements to keep them together Conduct electricity as a pure solid Polar No Non No Conduct electricity in aqueous solution Polar No Non No Dissolves in polar solvent (water) Polar Yes Non No Dissolves in non polar solvent Sometimes Yes

potential energy diagrams (LOOK AT GRAPH IN BOOK)

Electrons each have -1 charge Electrons are attracted by both nuclei The potential energy stabilization (called coulomb stabilization) is maximized when the electrons are bw the 2 nuclei since that minimizes R (R in denominator of eq so smaller that is bigger whole thing is) EQ in book Potential graph- as get closer will find position where they are most happy (D/r)- atomic radii might overlap a little to stabilize bond but not too close or will start to repel graph -- bond radius is distance from y axis to middle of u shape (where chemicals are at their most stable state)-- binding energy is the distance vertically from bottom of u shape to x axis

ionization energy

Energy needed to remove an electron from an atom Trends in ionization energies How much energy it requires to bump off one electron— fairy consistent up and down and left to right Alkali has least ant of energy and ones on right the most— they want to hold onto the electron the most As go right increasing effective nuclear charge As go down in table going up in valance shells As go down to bigger atoms get smaller bc those electrons are father from nucleus so less charge keeping them in (as up in valance shells down in charge that keeps them in nucleus) Cs is smallest- gives up electron easily, largest are noble gasses- requires a lot of energy to pull electron away from it

ionic solids v molecules and cations and anions

Ionic solids comprise of positively and negatively charged atoms. These are not discrete ions, instead a vast sea of charges holds the ions together in a crystal. Molecules are discrete groups of atoms bond by strong forces, pairs of electrons, making neutral bonds. Pos charged ions= cations Na+ sodium cation Ca2+ calcium cation An example of a "polyatomic" (many-atom) cation is the ammonium ion, NH4+ Neg charged ions = anions A negatively charged chlorine atom is an anion and is denoted CI-. An example of a polyatomic anion is the carbonate ion, CO32- ** Metals typically form cations and non metals typically form anions Metallic elements usually form cations by electron loss and non metals typically form anions by gaining electrons group 1 and 2 elements usually lose electrons to try to become like a noble gas -- if only have 1 valenece electron pretty willing to give it up **metals towards left side of table, metalloids (zig zag b to at) non metals on right -- non metals want bc have more VE-- farther over column wise

SI units-- cheat sheet??

Kilograms, meters, seconds, moles. Kelvin, ampere, candela A angstrom= 10to the -10 meters

mixtures vs solutions vs alloy

Mixtures >2 elements / compounds mixed together Can be separated by physical properties Dissolving in polar/ non polar solvent , magnet etc. , filtering Solutions <1nm Never settle out Single phase Solute can not be filtered out Alloy Homogenous mixture of 2 or more elements, one of which is a metal Usually mixing 2 metals to make stronger metal Alloying tin with copper to make bronze 1 has to be metal and the resulting material has metallic properties

molar mass v atomic mass and conversion how to w particles and moles

Molar mass- grams/ mol 1 mol of carbon has a mass of 12.011 g ( that is molar mass on periodic table) 2 mols of carbon- 24.022 g Atomic mass- relative mass relative to carbon 12 !! so molar mass is in mols and atomic mass is number relates to mols of carbon Molar mass is the mass of one mole - that's 6.022 x 10 - of atoms, or molecules. It must have a unit - typically, grams. When molar mass is given in grams, it has the same numerical value as relative atomic (or molecular) mass.!!! To convert bw moles and mass- use molar mass (from periodic table) To convert bw ,ass and particles- use both avogadros number and molar mass- two conversion factors Particles to moles and then to mass

VSEPR theory

Molecular geometry Specifies the positions of the atoms in terms of bond lengths and bond angles The shapes of molecules and their orientation critically affect their reactions, especially for biological reactions where molecules often must fit precisely into their receptors. A molecule adopts the shape that has the lowest potential energy (most stable). Molecular Geometry can be represented by different models Perspective drawing (not 3- d looking has lines and triangles) Ball and stick (3 d looking balls as elements connected w 3-d sticks) Space filling (bubbles. Connected to bubbles) Copy 3 MODELS of DRAWING Valence shell electron pair repulsion is. a simple but effective model for predicting molecular geometry VSEPR only applies to one atom at a time- usually the central atom Assumption (basis of VSEPR) A molecule minimizes the repulsive force between electron pairs. They will be as far apart on a sphere as possible to maintain the lowest PE. VSEPR theory- how to go about it Draw lewis structure and check formal charges Determine the steric number (SN) SN= # atoms bondied + #lone pairs (single double and triple bonds all just count as 1) Determine the electron pair geometry Determine the molecular geometry Want to make bonds farthest apart Molecular Geometries Start w election pair geometry Ignore lone pairs (only consider bonded atoms) Distortion from ideal geometry Non bonding electron pairs exert a greater repulsive force than bonding electron pairs Multiple bonds exert a greater repulsive force than single bonds

molecular mass

Molecular mass (relative) is determined by adding the atomic masses of all the atoms in the molecule (note- the atomic masses used are the weighted averages listed in the periodic table) So far water mass (H20)== 2 * mass of h + mass of ) Atomic mass is the mass of an atom and is given in a.m.u. (atomic mass unit). However, a molar mass is the mass of one mole atoms or molecules and is given in grams. one mole is equal to Avogadro's number ( NA=6.022×1023 ) atoms or molecules. all based on relative abundance- if more atoms have certain weight than others the atomic mass will be closer to their weight in the average-- one on periodic table

natural abundance and weighted avg atomic mass

Natural abundances Chlorine - 35 - 75.77 % Chlorine -37 — 24.23 % Average atomic mass- rarely use M is the relative atomic mass and add all isotopes masses together and then multiply it by the p which is their individual fractional abundances abundance of isotopes of a chemical element as naturally found on a planet. The relative atomic mass (a weighted average, weighted by mole-fraction abundance figures) of these isotopes is the atomic weight listed for the element in the periodic table. Weighted average atomic mass (75 % * 34.969 amu) + (24.23 % *36.966— average of isotopes of elements) = 35.453 amu (mass on periodic table) Mass on periodic table is added sum of protons and neutrons but average atomic mass is sum of masses of its isotopes each multiplied by its natural abundance (the decimal associated with percent of atoms of that element that are of a given isotope) Example Q What is the weighted average atomic mass of krypton that is composed of 3 isotopes 10% krypton- 79, 50% krypton-85, 40% krypton- 81 (.1*79) + (.5*85) + (.4*81) = 82.8 amu

isotopes

Neutral atom where number of neutrons is changed Unstable- radioactive and decaying Mass changes What can change with Chlorine, so it's still Chlorine and it's still electrically neutral? Number of neutrons

density

Rock sinks ship floats but float is more dense— that explains why- ship may have a lot of mass but the mass is spread out into a large enough volume that it is able to float whereas the rock is denser and floats Density decreases as temp increases Why- bc w change in temp you see change in volume (mass will not change) Density increases as pressure increases

precision v accuracy

Same answer every time when doing something- that is precise- not neccesarily middle of bullseye but all dots close to each other— usually results in systematic area Accurate— hitting middle of bullseye but kind of spread out in middle of bullseye- answers not neccesarily all close to eachother

Ion symbols and cations and anions

The chemical symbol (letter(s)) is determined by the number of protons. For example, the fluoride ion can be written as F- The F indicates that it has 9 protons The - indicates that it has one more electron than protons, so it has 10 electrons. Example prob How many electrons are in Mg2+? (answer= 10 electrons) Mg has 12 protons but the 2+ charge means that it has lost 2 electrons ***only thing that can change is neutrons (isotopes) or electrons (ions)— protons will always be same number Anions More electrons than protons Cations More protons Macroscopic matter is net neutral - neg charged ion must be paired w positive ions

molecular v complete ionic v net ionic eq

The molecular equation in electrolytic form (dissociated), is called the complete ionic equation: Ag+(aq) + NO3-(aq) + Na+(aq) + Cl-(aq) è AgCl(s) + Na+(aq) + NO3-(aq) (electrolytic form bc that is saying that when in solutions all these ions have broken up) We notice the same ions on each side of the equation, which means they play no role in our chemical reaction. These are called spectator ions. Rewriting our complete ionic equation without spectators, gives us the net ionic equation. Ag+(aq) + Cl-(aq)—- AgCl(s) (just including what is actually important in the creation of the precipitate) This net ionic equation is the net change taking place in the reaction when Ag+ ions combine with Cl- ions to form the precipitate solid silver chloride, AgCl. When put in solution- break up for eq and treat individual ions Cancel out ions that are not involved in actual reaction — if have same thing on both sides cancel them out to make reduced for of eq- then simplify eq— will be called NET ionic eq — comes from complete ionic eq *********Protons will react with hydroxides If weak acid that doesn't complexity dissociate you have to keep it in net ionic equation!!- going to keep the acetate together and in eq bc telling you you still have acid present in solution during neutralization reaction **weak acid or base have to worry about bc wont fully dissasociaste

percent yield and what values mean

Theoretical Yield - Maximum amount of product (mass) that could be formed from a reaction.- Calculated from a stoichiometry approach Actual Yield - Mass of product, measured in the lab. Percent Yield = actual yield/theoretical yield x 100 % How to maximize percent yield? Maximizing percent yield generally requires skill in choosing the best precursors, best conditions for the reaction, and using the best techniques.

Not all soluble ionic compounds dissociate!! Common weak electrolytes

They are soluble They do react But they do not fully break apart into ions Weak acids and bases are ex of above Weak acid that is electrolyte but does not dissociate HC2H3O2 Weak base that is electrolyte but does not dissociate NH3 Weak salt that is electrolyte but does not dissociate HgCl2

Atomic radius

This depends on how you measure it-will have diff results if measure in gas stage vs crystal phase - fuzzy boundary Atomic radius in solid is well defined bc can measure radii of crystals Atomic radius increases as go down a group to larger atoms also increases if go RIGHT to LEFT column wise (opp of what would thing) Across a period, atomic radii get smaller due to increasing effective positive charge on the nucleus Restraints- nucleus is pos and electron is neg so they are attracted to eachother- counter balancing you have all these electrons w neg charge- if you pull them to close they will be too restrained so repulsive forces will pull them apart As go from left to right adding proton (increasing charge 1 from nucleus) but also adding electron— problem is electrons are delocalized so in net addition of given electron does not completely cancel out charge- so have increase in effective pos share in nucleus!! Greater it is more it will want to associate w electrons Down a group, atomic radii get larger due to electrons in higher energy levels (shells)— bc filling more shells - of course will get bigger 89 Going down has bigger affect than across The radius of a sphere measured as the distance between nuclei in solids

classifying compounds by the type of element

Two nonmetals are molecular: Water (H2O) is an example of a binary molecular compound (non metal w non metal) A metal and a nonmetal are ionic: sodium chloride (NaCI) is an example of an ionic compound. (cation with anion) Polyatomic ions consist of three or more atoms. The ions can have either a positive or negative charge. These ions are molecules which then can form a part of an ionic compound

A sample of liquid N2H4 (hydrazine) is decomposed to give gaseous N2 and gaseous H2. —-N2H4 (l) → N2(g) + 2 H2(g). The two product gases are separated and the nitrogen occupies 15 ml at room temperature and atmospheric pressure. What volume of hydrogen will be produced under these same conditions?

answer= 30 ml bc for every n2 there are 2 h2

Electron affinity

is defined as the energy released when an electron is added to an atom. The reaction X(g) + e-èX- is called an electron attachment reaction. • X- is called an anion. • If energy is released, the anion is stable. • If energy is required, the anion is not stable. • The largest electron affinities are for F and Cl. Both electron affinity and ionization energy can be measured in the laboratory. Energy given off when an extra electron is added to an atoms How much energy is released if we put an electron on an atom (have some that require energy to do that but majority actual want and will take on that electron)!!!*** Opp of the ionization energy stuff Negative values: energy is required to add an electron Pos values: energy released or favorable General trends are same as disassociation

electron configurations

the electron configuration of each element is critical for its observed chem properties S- block, f- block etc are all indicating electronic config— they have orbitals and are picky about how they assemble— each line on periodic table responds to corresponding shell— as build up putting electrons on and finishing shell- when have complete shell you have noble gas- they are happy and stable Then have orbitals (s orbitals- only space for 2 electerons) p- only 6 d- space for 10 F- space for 14 Electrons go to diff spots based on reactivity Predictable way in which we can put in electrons Electrons have spin- will assemble so spin is app and will never have more than 2 electrons in 1 orbital - if there is space for electron to take own orbital it will — some some orbitals will only be half filled and then if need be will go back and start filling with second

trends in bond length

triple bond has more bonding electrons than double bond and single bond- more bond energy bond length decreases as bond strength increases!! so inverse relationship 2 above longer bond length bw bigger atoms • Among members in a group in the periodic table, larger atoms have longer bond lengths. - Which would have the longer bond length? HCl or HI • Bond lengths (of the same type) between a given pair of atoms does not change much in different molecules. For some pairs, there are 2 types of bonds (C-C for example can have one 2 or 3 bonds) Bond length and bond energy are fairly uniform Bond lengths bw 2 elements will usually have charectrositic distribution of lengths- C-C vs C-H each will have their own individual lengths As go up in number of bonds length gets steadily shorter!!!***

reduction oxidation reaction 2Mg +O2—-2 MgO ex

• Oxidation:Loss of electrons ,or gain of Oxygen atoms • Reduction:Gain of electrons,or gain of Hydrogen atoms • Oxidation numbers increase with oxidation,decrease with reduction (bc becoming more neg!!) • Reducing agents reduce the species they react with • Oxidizing agents oxidize the species they react with Reduction Oxidation reactions range from common: combustion, corrosion, to elaborate: photosynthesis, metabolism and metal extraction reactions Here Mg atoms lose electrons to form Mg2+ ions, and the oxygen O2, gains electrons to form O2-.

oxidation state trends in periodic table

*****************************************electronegativity increases left to right but then decreases w as go up atomic number in group so as go down columns Starting w alkali metals all the way to the left (+1 charge)— they don't mind giving away their one valance electron- they are not too electronegative — they are good candidates for forming ionic bonds These are usually what is being oxidized- they give away an electron **bc alkali metals willing to give electron their oxidation state is +1 (giving an electron away will make atom more pos) One over- alkaline earth metals not too electronegative so they could give away 2 valance electrons and be oxidized — in a typical ionic bonding situation they would be more likely to give 2 electrons — bc it would take a lot for them to complete their valance shell Group 7- halogens- very electronegative sitting on right side of periodic table - one electron away from being satisfied- so usually reduced (have an oxidation state of -1 ) (will want to gain an electron) so oxdation state is -1-- means they are in a state in which they want 1 more electron to be satisfied Group 6- where O sits- (oxidizing is doing to something what oxygen would have done so oxygen is a very good oxidizing agent — will take electrons— will take 2 electrons so oxidation state is -2)

4 common gas forming reactions (CHEAT SHEET???)

1) H2CO3(aq)—-CO2(g) + H2O(l) 2) H2SO3(aq)—-SO2(g) + H2O(l 3) NH4OH(aq)—-NH3(g) + H2O(l) 4) Mmetal + HCl—-H2(g) + MmetalCl

The words "oxidation number" and "oxidation state" are interchangeable. Assigning oxidation number to elements use simple rules: and common oxidation charges-- CHEAT SHEET??

1) The oxidation numbers of uncombined elements are zero (0) (free elements or atoms with themselves) i.e. H2, O2 F2, Cl2, Li(s)- U(s). 2) The sum of the oxidation numbers of all the atoms in a species is equal to its total charge. Common oxidation numbers Hydrogen is H+ unless it's with a metal when it becomes H- • The reason for this is that most metals are more electropositive than H so H takes on the negative charge. • Groups 1 and 2 oxidation numbers are equal to their group number. • Halogens are -1 unless the halogen is in combination with oxygen or another halogen higher in the group. E.g. ClF2 • Oxygen is O2- except when combined with fluorine. Less common are peroxides (O22-), superoxides (O2-), and ozonides (O3-).

first ionization energy rules

1. Across a period, 1st Ionization energy increases due to increasing positive charge in the nucleus (more pos charge in nucleus more the electrons are going to want to stay) 2. Down a group, 1st ionization energy decreases because the electron is in a higher energy level, further from the nucleus and more shielded by other electrons 3. Effect of electron pairing after half-filled p-subshells is seen, decreasing 1st I.E. for 4th electron Oxygen breaks trend going right to left bc if loses electron it is actually happy - there are lots of anomalies

first type of lewis structure exception (WRITE ON PAGE THAT N CAN TAKE EXTRA 1 ELECTRON)

1. Odd electron molecules are known as (radicals)!!!1. Most are not stable for more than a few microseconds- but some are stable over long periods such as NO — if election left over have to figure out where to put it— so look at formal charges and then may add one over the 8 rule to carry the extra electron

Which of the following chemical equations refers to the second ionization of Ca? A. Ca(s) + 2e- —-Ca2-(s) B. Ca(g) —- Ca2+(g) + 2e- C. Ca(s)—-Ca+(s) + e- D. Ca+(g)—-Ca2+(g) + e-

1st ionization would be Ca (g) —- Ca+ (g) + e- (took 1 electron away gave it charge of pos one) 2nd ionization would be Ca+ (g)—- Ca 2+ (g) +e-

neutralization reactions

Acid base reaction is a Neutralization reaction Neutralization reactions take place between a strong acid and metal hydroxide: Acid + metal hydroxide salt + water The net outcome of any strong acid base neutralization reaction is the formation of water. !!!!!

acids and bases and diff definitions

Acids and bases change the color of certain dyes known as indicators Acids (tart- turn red) Bases- bitter (turn blue) Svante Arrenius- swedish - in 1884 said An acid is a compound that contains hydrogen and reacts with water to form hydrogen ions (H+) HCl (aq)—- Cl- + H+ base is a compound that produces hydroxide ions in water (this is not wrong but just limited) NH3+ H20—- NH4+ + OH- Limitations to the Arrehnius def 1. It is specific to one particular solvent, water. 2. Not all base reactions produce hydroxide ion, OH-. 3. The key process in an acid and base reaction is a proton (H+) transfer (little to do with OH-). In 1923, Thomas Lowry in England and Johannes Brønsted in Denmark, came up with a proton (H+) transfer idea. An acid is a proton donor and a base is a proton acceptor Ex: HCl+ H2O—- Cl- + H3O+ HCl is the proton donor and H2O is the proton acceptor HCI releases a hydrogen ion, H+, to water, producing hydronium ions (H3O+ ) and chloride ions. H2O accepts the hydrogen ion to form H3O+, water is acting as a H2O accepts hydrogen to form H3O+ water is acting as a bronzed base in this reaction Classifying Acids and bases Bronzed lowry acids and bases further categorized based on extent of deprotonation or protonation Strong acid is completely deprotonated in solution Weak acid is incompletely deprotonated in solution A strong base is completely protonated in solution A weak base is incompletely protonated in solution strong 100 ionization

covalent bond and what bond order is and diff number of bonds

All of bio run by covalent bonding (C,O,N and then a couple others— all of those are non metals) Chemical bond formed bw atoms by the sharing of electrons and overlapping electron clouds (usually between two non-metals) What makes them form covalent bond? Want to have noble gas like outer shell- want to have 8 valence electrons- want to hold onto electrons bc non metals so will share electrons to fulfill that Double and triple bonds Atoms can share four electrons to form double or 6 to form triple Number of electron pairs is the bond order O2 - 2 of these can share 2 each w each other Electronegativity Polarity refers to a seperation of positive and neg charge. In a non polar bond the bonding electrons are shared equally In a polar bond electrons are shared unequally— polarizes towards more electronegative Degree to which they are polarized is dependent on relative electronegativities of atoms bonded together As go across columns counting up to 8 in valence electrons!! — those are ones population outer shell P and S orbitals

moles

Avogadros number- number of atoms in 12 g of 12C and has symbol Nsubscript A Na= 6.0221413*10^23 There are 6.0221413*10^23 molecules in a mole The relative masses on the periodic table are also the molar mass (grams per mole) for the atoms Molar mass corresponds to grams per mole for that substance Best experiments use x rays to measure the distance bw atoms in a crystal Moles have different weights but there are the same number of molecules in each mol of something — why we do this bc gets to standardized account in terms of number of atoms

Rank Ba, Mg, and Ca in order of increasing first ionization energy.

Ba<Ca<Mg As get down to bigger atoms electrons farther apart so easier to pull them out (as go down rows easier to pull away)

compounds- metals, non metals and metalloids

Can be broken down through chem reactions or change More than one atom connected together 75% of elements are metals — can conduct electricity , malleable, ductile Non metals- insulators (do not readily allow heat or electricity to pass through- used to protect from electricity flowing through things) Metalloids- can act as a metal or a conductor

element

Can not be broken down into smaller constituent parts through chem reactions (chem change) Unique atoms Atoms are not connected Occur as single atoms except N, O , F, Cl, Br, I, Bh (diatomic elements — always have 2 under them) Some elements have similar microscopic properties Arranged in the Periodic Table during 1800s according to observed chemical properties. Chemical Reactivity Ionization Energy Electronegativity Number of Valence Electrons Oxidation Number atoms= particulate view ***** Molecules w 2 atoms are called diatomic- if has more than 2 then called polyatomic

changes in physical vs chemical state and indicators of chem change

Changes in physical state or chem state Physical Measured w out changing subs identity Chem Measured by changing chem identity Only measured through the course of a chem reaction Physical change (affects only physical properties) and chem change (a new substance is formed) Indicators of chem changed Heat or light is given off or taken in Color change Precipitate is formed (solid) Gas is given off (not due to change of state) Physical changes Trans bw solid liquid and gas (phase changes) Physical properties Changes How do we know going on? Physical- affects only physical prop Chem- new subs is formed

chemical bonding ionic vs metallic bonds What happens to valence electrons type of compound structure types of atoms electroneg diff

Chemical Bonding- Ionic Ionic compounds consist of a crystal lattice of positive and negative ions Unit cell: simplest repeating unit in a crystal lattice Metallic bond Chemical bond formed by the attraction bw positive metal ions and bands of free electrons Metal and more metals What happens to valence electrons Metallic majes a sea of electrons- delocalized Ionic Electron transfer type of compound structure Metallic Crystal lattice Ionic Crystal lattice Types of atoms Metallic Metal and metal Ionic Metal and non metal Metal and non metal Electronegativy diff Metallic Doesn't matter Ionic >1.7 then becomes more appropriate to consider atom as ionic —- if electro diff is more than this

molecular v condensed v structural formula v Electrostatic potential surface

Chemical formulas represent the composition of elements The subscripts show the relative number of atoms in the smallest unit Estrone and testosterone- male and female hormones only differ by a few atoms Molecular vs. condensed formulas Using methanol for an example Molecular formula = CH4O Condensed structural formula= CH3OH— this way indicates atom arrangement Structural formulas indicate how atoms are linked together, lines represent chemical bonds (this is more picture way w lines and letters coming off) Electrostatic potential surface (elpot) — these show electric charge distribution Shows 3-D structure of a molecule w huge bubble around it and red tint indicates negative potential due to negatively charged electrons (probably free electrons in that area of molecule (like on outside where there are lone electrons not bound to another element of molecule) Blue tint indicates positive potential due to a positively charged nuclei

compounds and organic v inorganic

Compounds are electrically neutral consisting of two or more diff elements Compounds are classified as either organic or inorganic. Organic compounds contain carbon and usually hydrogen. There are millions of organic compounds, including fuels, sugars, and most medicines. Inorganic compounds are all the other compounds; they include water, calcium, ammonia, silica, hydrochloric acid, and many more.

Democritus, Lavoisier and Dalton

Democritus Developed the word and concept of an atom. Piece of an element that retains its properties Lavoisier French chemist- Law of conservation of mass Law is in chem reaction matter is neither created or destroyed- just transferred Dalton 1808- atomic theory Each element made up of atoms and they are indivisible Atoms of one element are the same (properties) but differ from other elements Atoms can combine in simple ratios into compounds Cannot be created or destroyed in any chemical reaction Given compound has same relative number and kinds of atoms

dissolving v dissasociation and soluble v insoluble

Dissolving vs dissociation- sugar dissolves in water but does not dissociate into ions- can dissolve vodka in water but it is a nonelectrolyte- oxygen dissolves in your blood but it does not ionize Soluble vs insoluble Soluble substances dissolve in a solvent K2Cr2O7 is soluble in water; sugar, vodka and oxygen also dissolve (are soluble) Soluble implies water (aqueous) as solvent unless otherwise stated Soluble ionic compounds will both dissolve and dissasociate Soluble molecular compounds will dissolve and may or may not dissociate Insoluble substances do not dissolve significantly in a solvent. A copper wire is insoluble in water, like rubber or skin.

metallic bonding and trends in metallic character

Electrons are mobile. This enables metals to conduct an electric current. Also accounts for their luster and malleability being deformed w out shattering Electrons will delocalize into sea of electrons surrounding metal nuclei- must remain charge of neutrality- put number of neg on one side but then same amount has to be coming out other side to be constant Metals dont want to grab electrons that much so electrons way more likely to delocalize Trends in metallic character —- upper right is least metallic bottom left most metallic Trends Metallic character: how loosely electrons are held Opposite trend from ionization energy (attraction for electrons it already has) Opposite trend from electronegativiy (attraction for bonding electrons) ** Metallic character Ability to lose a valance electron

empirical form and molecular form and steps needed to find emp form

Empirical Formula The smallest whole-number ratio of atoms in the compound. Molecular Formula The actual number of atoms in the molecule. A whole number multiple of the empirical formula for that compound. Example Empirical Formula = CH2 = Molecular Formula = C2H4 Ethene Molecular Formula = C3H6 Propene Molecular Formula = C4H8 Butene Steps needed to find empirical formula 1. Find number of moles for each element 2. Divide all results by smallest number of moles present 3. Find smallest multiplication factor needed to make all numbers whole 4. Multiply all formula subscripts by this factor

extensive v intensive

Extensive vs. intensive Extensive depends on amount of matter - doesn't have anything to do w substance itself (volume, length) doesn't relate to identity of matter itself Intensive dont depend on amount of matter but are inherent to the matter itself Can be used to identify a substance Boiling point, flammability **ratio of 2 extensive properties often results in an intensive property (density = ratio of mass to volume so density is intensive.) -density intensive derived from to extensive properties- it is this because density of water for example changes at different temperatures

suspension and colloids

Gross mixture Suspension Particles > 1 μm, settle more rapidly Solute can be filtered out ex: muddy water, milk, heterogenous Colloid Particles 1 nm - 1 μm diameter Settle slowly Solute can not be filtered out a homogeneous, noncrystalline substance consisting of large molecules or ultramicroscopic particles of one substance dispersed through a second substance. Colloids include gels, sols, and emulsions; the particles do not settle and cannot be separated out by ordinary filtering or centrifuging like those in a suspension. Gel like substance ex: dust, fog, jello

Faraday and compounds 1839

H2=H20=Ar Atomic structure is related to electricity Ben Franklin had discovered earlier that there are 2 types of charge pos and neg His discovery leads to idea that atoms have an electric component — electrical forces were responsible for the joining of atoms in compounds Compounds Elements in fixed, small whole # ratios Electrically neutral atom- # electrons= # protons

homogenous solution

Homogenous Solution - Particles < 1 nm (atoms, ions, molecules) - Never settle out - Single phase - Solute cannot be filtered out ex: Saltwater, soda, kool aid Solvent > 50% and solute < 50% ***Mixtures of two or more gases are ALWAYS homogeneous and ALWAYS solutions.

acid naming

Hydro- add name- ic (ex: hydrochloric acid) If has ate or ite in name, must be acid — most likely add O Complex acid compounds have oxygen in them. For an acid with a polyatomic ion, the suffix "-ate" from the ion is replaced with "-ic." Polyatomic ions with one extra oxygen (as compared to the typical polyatomic ion) have the prefix "per-" and the suffix "-ic." Polyatomic ions with one fewer oxygen have the suffix "-ous"; ions with two fewer have the prefix "hypo-" and the suffix "-ous." H+ non metal or H + polyatomic ion If H+ non metal hydro+______ ic ex:HCl=hydrochloric acid HF= hydrofluoric acid If H + polyatomic ion ate—- ic and ite—-ous HClO4= perchloric acid HClO3= chloric acid HClO2= chloric acid *** HCN= hydrocyanic acid **** HOCN= Cyanic Acid

different options in terms of electrolyte if If soluble (Aq)- dissolved vs if insoluble

If soluble (Aq)- dissolved Could only dissolve Would be non electrolyte Sugar, vodka, oxygen Could partially disassociate Would be weak electrolyte Acetic acid, most organic acids Could disassociate Strong electrolyte KMnO4 and many more If insoluble Do not dissolve Nonelectrolyte skin, copper wires, tires

Dilutions

If you dilute the solution, the number of moles does not change. So in the diluted solution, the product of the molarity and volume must be equal to MAVA. This is generally written as MAVA = MBVB Also often need to be able to dilute aqueous solutions to a different (lower) concentration Stock solutions- solutions that are concentrated and then you will dilute it to some concentration (need to figure out how to add enough water to create buffer you will work with) number of moles in a solution A can be MaVa (mol/L * L)= mol! Molarity times the volume of the solution will give you the number of moles in that solution!! If you dilute the solution, product of the molarity and volume must be equal to MaVa Generally written as MaVa= MbVb—- bc moles wont change !!— just diluting it

precipitation reaction

In a precipitation reaction, (aq) indicates a substances dissolves in water and (s) indicates a solid precipitated 2 solubles to start— then insoluble comes out (precipitate) Soluble compounds can be strong, weak and nonelectrolytes; these may or may not disassociate. Insoluble compounds are nonelectrolytes, and do not disassociate.

limiting reactant how to do

LM is the one that runs out first Excess reactants are ones that don't run out Approach Convert reactant masses to moles and determine the theoretical yield in moles from each reactant (how much you would expect) One that corresponds to production of least product is your limiting reactant

Law of constant compostion, law of definite proportion, law of multiple proportion

Law of constant composition The composition of a substance is always the same, regardless of how the substance was made or where the substance is found Example: 1 molecule of water (H2O) is always (and only) made up of 2 atoms of hydrogen and one atom of oxygen. Law of definite proportion The same compound always contains the same elements in the same proportions by mass or percentage of the total. Example: H2O contains 1 O and 2 H atoms H2O2 contains 2 O and 2 H atoms Law of multiple proportion Whenever the same two elements form more than one compound, the mass ratios are small whole number multiples.

Lewis structures and rules for making them

Lewis structures for covalent molecules— predict what covalently molecule will look like Only valence electrons are shown Each atom should be surrounded by 8 electrons (except hydro and helium- both form duets) Shared pairs are written as H;H or H-H Unshared pairs are called lone pairs and they don't contribute to bonding- called non bonding electrons Rules for making them Predict the central atom The single element in the formula (if have one S and a bunch of O— then S is central) Carbon can make 4 bonds so usually in middle Element w the lowest electronegativity will be in middle!!! Count total number of valence electrons (adsd up for all atoms together) Count the number needed to fill all the octets on those atoms The number of bonds #B= (#Octet atoms have-#VE)/2 Put electrons around atoms to make octets If there are not enough electrons, use double bonds/triple or ring Use formal charges to choose the best option- asses whether realistic structure or not (calculate formal charges on structure)

why masses on periodic table have no unit and mass number what it is How many neutrons are in 1000 atoms of Cl-37 Cl 37 is an isotope that has 20 neutrons

Mass of a single atom of carbon (12C) is set at exactly 12 u where u is an atomic mass unit (amu)— mass of other atoms is measured relative to this ** masses of the periodic table have no unit bc they are related to 1/12 the mass of C12— all average atomic masses Mass number= number of protons + number of neutrons Cl 37 is an isotope that has 20 neutrons and 17 protons so answer= 20000 **** isotopes the dash and then number denotes the mass number and then subtract the number of protons in a regular nitrogen to get neutrons bc isotope is change in neutron Element symbol Top left corner = mass number= protons +neutrons Bottom left corner= number of protons

potential energy diagram for ionic bonding LOOK AT PIC IN NOTES

Na and Cl bonding (they both have charge which took energy to happen but we are not worried about that now) X axis- internuclear distance bw ions- if far apart don't know each other and then push them closer and at some point if push them too close those electrons are squished up again so then there is a repulsion again after the attraction- there is a balance End up w middle ground- min potential energy is equilibrium bond distance — bc solid it would be sum of ionic radii bw 2 ions Energy from r up to 0 is dissociation energy- will need to put in energy to break that bond apart Y axis potential energy

naming ionic

Name cation first (metal or polyatomic cation) Name anion second (non metal) and then change suffix to ide For transition metals must include roman numeral after them so that you know based on other element that they are paired w, what their charge is for this specific bond

powers of ten-- cheat sheet???

Nano- 10^-9 Kilo- 10^3 Centi- 10^-2 Micro- 10^-6 milli- 10^-3 Mega- 10^6

Nucleus, proton v electron and valley of stability

Nucleus is 1 mm (barely larger than period at end of sentence) compared to atom which is 100 yd stadium ** nucleus is most of atoms mass but very little amount of its size Proton 1.672 * 10- 24 grams The nuclear charge (charge of the nucleus) is always the same as the number of protons. Electron .00091095* 10 -24 g WAY smaller than proton Valley of stability" 1/1836 of proton N/P ratio To see which is more stable you look at element and do mass number- protons / protons and the lower number is more stable when comparing 2 elements!! is a characterization of the stability of nuclides to radioactivity based on their binding energy.[1] Nuclides are composed of protons and neutrons.

balancing chem reactions and hints

Number and identity of atoms in reactants must be the same as products (balanced reaction- cant create or destroy atoms) Number of molecules are called the stoichiometric coefficients Need to add coefficients and subscripts - write out what have initially of each from subscripts and then add coefficients — can not change subscript to balance equation — changing identity of compound then Hints for balancing equations Look first for elements that appear only once on each side of eq Create polyatomic ions as groups (they are stable so don't worry about breaking them apart) Leave lone elements until end

pure substance vs mixture

Pure substance Element (goes through chem change) and becomes Compound Mixture (homogenous has same uniform appearance throughout— many referred to as solutions and heterogenous mixture consists of visibly different substances or phases

Rules for naming molecular compounds

Remove ending of second element and add ide Prefixes are used to dictate number of element present in given compound (mono, di , tri, tetra, penta, hexa, hepta, onto, nona, deca) If only one of first element- drop prefix Parenthesis used to tell charge of something that could have many charges- what is in parenthesis depends on element it is paired w- to make charges equal and opposite

multiple ionization

Removing an electron from an inner shell (core) require even higher energies Subsequent ionizations are always higher in energy than previous ionizations- greater Columbic imbalance More electrons you take off more energy needed to take them off

resonance forms (and example of what bond order is w O3) and formal charge

Resonance forms are written when there are 2 or more equivalent lewis structures (same atoms in same arrangement but maybe double bounds in diff places- that is resonance form) way of describing molecule that may have mixture of single and double bond present for O3 the double bond can be on one oxygen or the other and they create structures w the same formal charges so for bond order you take the average of the values in all the structures (both structures create 3 bonds so you take 3/2 (bc 2 structures= 1.5) Formal charges Talking about ownership Want to calc diff bw ownership and how many it had when you consider it in the beginnning- how many did it own in final compound — how many does it own in final compound vs how many did it have prior to bond Ideal case for lewis structure is that formal charge is 0— hasn't changed!! from one they were just an atom alone If + or_ that means deviation bw what it had alone vs when bonded Are based on the number of electrons assuming all are shared EQUALLY Overall the formal charges add to result in the charge on the ion. • The only purpose of formal charges is to determine the preferred Lewis diagram. • If 2 or more Lewis diagrams can be formed from the same elements, the one with the lowest formal charges is best. Formal charges >1 are not found in good Lewis structures.

Geometries for the distribution of electron pairs for different steric numbers— not molecular geometries they are the first step in determining molecular geometries!! WRITE ON CHEAT SHEET

SN=2 (linear electron pairs) Can only be linear SN=3 (triginal planar electron pairs) Trigonal planar or bent depending on how many lone pairs vs bonds you have Trigonal planar If there are no lone pairs on the central atom (angle= 120 if all outer atoms are the same (BF3)) Bent - if 1 electron pair with an angle < 120 (H-N=O) SN=4 (tetrahedral electron pairs) tetrahedral- if no lone pairs (CH4) Pyramidal (or trigonal pyramidal) if 1 lone pair (NH3) angle is smaller than tetrahedral angle of `109.5 Bent- if 2 lone pairs(H2O) angle is smaller than tetrahedral angle Very polar H2O Once get shape want to figure out how polar molecule is SN=5 Trigonal bipyramidal electron pairs There are 2 axial positions and 3 equatorial (PF5) Each axial has 1 repulsion that is @ 180 and 3 repulsions @ 90 Each equatorial has 2 @90 and 2 @ 120 Lone pairs fo on equatorial positions to minimize repulsion 1 lone pair= seesaw (SF4) 2 lone pairs= distorted T (CIF3) 3 lone pairs= linear (XeF2) SN=6 Octahedral electron pairs Octahedral is no lone pairs (SF6) Square pyramidal if 1 lone pair (IF5) Square planar if 2 lone pair (XeF4)

What happens at the particulate level when salt dissolves into water? (NACl)

Sodium cations and chlorine anions separate and float around in solution Water is not just the background— chloride and sodium ions are interacting w them in weak manner- not bonded but balancing some of the charges (kind of floating around them)— hydrogen has part pos and O has part neg so Cl ion which is neg will have protons from water congregate around it For sodium cation oxygen of water will go towards cation and this will make that ion happier to be in solution Water stabilizing compound when it breaks into ions When put aquous next to something- really is another species bc putting insulation is efficient way of making atoms accessible for reaction

second type of lewis structure exception (SEE ABT PUTTING BR ON PAGE)

Some group 2A and most group 3A elements form molecules where they are octet deficient ( the groups are columns so 2 A is molecules in column 2 then skip transitions and 3A is on other side) Br-Be-Br Have small number of electrons to start out w- cant achieve a bonding scheme to give them an octet Most elements in group IIA form ionic bonds so the octet rule applies. But Be forms polar covalent bonds with some of the halogens Br does not form double bonds Not enough electrons to provide all molecules with an octet with single bonds If have good formal charges then still have good structure— so can always check w that!!! ***SO USUALLY HAPPENS BC BR DOESNT FORM DOUBLE BONDS AND CANT FILL OCTET W SINGLE BONDS

Third lewis structure exception rules for how to do them

Sometimes occur in compounds with group 5A,6A,7A, and 8A atoms as central atom Example SF8 #VE= 6+6(7)=48 #O=7(8)=56 #B=4 (can see write here that you cant make a model w just this amount of bonds)- minimum of 6 bonds are needed So the electrons on the central atom are expanded to have 12 electrons in 6 bonds. That leaves 48-12=36 to distribute around the 6 fluorine atoms, giving all elements formal charges of 0. Make minimum amount of bonds if you know the 6 bonds w sulfure exist (sulfur now has 12 electrons around it- unused D orbitals is what enables sulfur to handle those bonds rules for how to do them For Valence shell expansion Give single/double bonds to all the terminal atoms and add electron paire to the terminal atoms to give them octets If there are any left over electrons- they go as non bonding pairs to the central atom This preserves the octet rule for the terminal atoms and allows the central atom to have its valence shell expanded

strong v weak acids

Strong vs weak acids You will never see HCl in water (bc strong acid) but you will see a lot of acidic acid that is propriated in water —-because it completely dissociates (100%) into H+ and Cl- ions. On the other hand, for a weak acid like acetic acid, CH3COOH, you will see mainly the acid and only about 1% acetate, CH3COO- . Only small fraction of acids actually go into water and this is important

5 types of reactions

Synthesis- put things together A+B—- A-B 2N2 + 5 O2—- 2N2O5 Decomp Break things apart A-B——A+B 2HgO—- 2Hg + O2 Combustion CxHy +O2—- CO2 +H20 (what we do when burn— reacting something w Oxygen (oxidation) C7H16 + 11O2—— 7 CO2 + 8 H2O Single displacement A-B +C—— A-C + B Swapping anions CuCl2 + Mg —— MgCl2 + Cu , double , A-B + C- D—— A-C + B- D 2 NaOH + CaBr2—- Ca (OH)2 +2NaBr

laws of combining volumes

The ratio of the volumes of any pair of gases in a gas phase chemical reaction (at the same temperature and pressure) is the ratio of simple integers Ratio of volumes of gases in gas phase reaction is the ratio of simple molecules (same temp and pressure) Example: 2 volumes of hydrogen + 1 volume of oxygen give 2 volumes of water vapor The ratio between the volumes of the reactant gases and the gaseous products can be expressed in simple whole numbers. 2 molecules of hydrogen + 1 molecule of oxygen = 2 molecules of water. The law of combining gases was made public by Joseph Louis Gay-Lussac in 1808.

Polar molecules and dipole moment and biochemistry

The total dipole moment of the molecule is the sum of the polarity vectors for each of the bonds Polar molecules have much stronger electrostatic interactions than non polar molecules. This affects their physical and chemical properties. Calculare electronegtaivity difference for all bonds in the molecule If all bonds are non polar- molecule is non polar If polarity is symmetric- molecule is non polar If polarity is asymmetric- molecule is polar Dipole Moment and Biochemistry Changing polarity of molecules can change pharmacological activity due to the principle that "like dissolves like" For example adding non polar groups Increases solubility in fat Facilitates passage of a drug into the bloodstream through the intestines Large biochemical molecules such as proteins fold because exposed groups have opposite polarity For example, an -NH2 (amide) group will attract a -C=O (carboxyl) group and the protein will fold to bring them closer together. The larger the electronegativity difference, the stronger the attraction will be toward a group of opposite sign. Pic of TIM barrel protein: The complex shape of this protein is due to the intramolecular interactions due to electronegativity difference bw the atoms (and the resulting dipole moments)

electrolytes-- put couple ex of each electrolyte type on cheat sheet?????

When in water some ionic or molecular compounds conduct electricity, some strongly and some not at all Distinguishing between electrolytes is how well they conduct electricity (so can dissolve but not be electrolyte-- it is more about how well they conduct electricty- produce cations and anions) non electrolyte Does not form ions in solution Can be solids or dissolve yet no ions are present Molecules Do not disassociate- remain intact weak electrolyte Barely ionize in solution- they mostly remain intact Acetic acid is a weak electrolyte Only about 1% will separate into proton and then those actually associate w water to make hydrogen ion Only a small fraction (1%) of CH3COOH molecules separate (slightly or dissociates negligibly) into hydrogen ions, H+, and acetate ions, CH3CO2- . Slight dissociation Complicated in these cases bc how much exactly will dissociate?- there will be different degrees of the weak ones strong electrolyte Only solutions w ions can conduct electrical current since ions have charge (so become charge carriers)— so compounds that dissociate w in solution are electrolytes — notice ions become free to move when the solid dissolves- electrolyte is just ion that is used to conduct electricity due to fact that it is mobile in solution

electronegativity

a measure of how strongly an atom attracts bonding electrons. Mulliken (mulliken defined this as average of electron affinity and ionization energy) and Pauling (grounded on measurements of dissociation energies- ant of energy takes to break molecules apart) scales (pauling is more grounded in experiments - what used in textbook) Generally, trends like ionization energy - how willing atom is to take an electron Across a period, electronegativity increases due to increasing effective positive charge on the nucleus. Down a group, electronegativity decreases because the bonding electrons are in a higher energy level, further from the nucleus and more shielded by other electrons. *All electronegative are pos- will take electrons Electronegativities are not assigned to the noble gases Larger electronegativity the more non metallic the atom Most electronegative atom is fluorine alkalis have lowest electronegs and are called electropositive. (all electronegativity values are positive). - The smaller the electronegativity, the more metallic the atom. - The lowest electronegativity is for francium and cesium (0.79 for cesium). ** Large electronegative means atom will NOT readily lose electrons

percent error calculation

calculation is the experimental or measured value minus the known or theoretical value, divided by the theoretical value and multiplied by 100%

1811- avogadros hypothesis

gas has same number of molecules Equal volumes at same T and P states that two samples of gas of equal volume, at the same temperature and pressure, contain the same number of molecules. Allows chemists to predict behavior of ideal gases Equal volumes of different gases (at the same T, P) contain equal numbers of molecules.

oxidation number

how many electrons are lost Oxidation is loss reduction is gain (reducing the number of the charge— going from charge of +5 to +3- that is a reduction )— where one thing loses electrons somewhere else gains in order to have balanced eq Oxidation number: Keeping track of electrons Free atoms have oxidation numbers of 0 Mostly worried about what we do when we have compounds Need to take into account total charge and then where those charges belong on compound — look at one element you know charge for and then deduce other one based on that Start w elements that almost always have same oxidation number and then go from there

diatomic elements-- cheat sheet???

hydrogen, nitrogen, oxygen, fluorine, chlorine,iodine, bromine.

3 ways of looking at chem

macroscopic view (match burning) Particulate view Talking about atoms and molecules and what is happening Chem Equation view Gives concise desc of what is happening at micro level and describes what see at macro level

equitorial atoms

make a straight line through center- center molecule and 2 others in line dont highlight - highlight others

ionic compounds and formal units

nic compounds form into crystals due to the vast sea of charges between anions and cations. Formal units are discrete, electrically neutral, units in a crystal. NaCI is the formula for the crystal containing one Na+ ion for each Cl- ion. The crystal for the binary compound, CaCI2, is formed from Ca2+ and Cl- ions in the ratio 1:2. 16

ionic bonds polar covalent bonds and covalent bonds and octet rule

onic bonds- electrons are transferred. They form between atoms with very different electronegativotes Chemical bond formed by the transfer of electrons electrostatic attraction of resulting ions (cation and anion) - The cation is typically a metal - The anion is typically a non-metal Polar covalent bonds- electrons are shared but not equally by atoms w intermediate electronegativity differences Covalent bonds- electrons are shared. They form between atoms with small or no differences in electronegativity. Octet rule atoms gain, lose or share electrons to have eight valence electrons (in their outer shell) Like to become like nearest noble gas - atoms gain, lose or share electrons to have eight valence electrons in their outer shell Refers to outer shell s and p orbitals— all will want to take 8 except H and helium Oxidation number- equals the charge on the ion

trends in bond energy

— dissasociation energy Depends on the size of the atoms, electroneg difference, and structure of molecule Bigger atoms not able to get as close together - bond is weaker Bonds generally grow weaker with increasing atomic number. - Which would have the stronger bond, HCl or HI? • Bond energies are generally within 10% in compounds with the same atoms. - For example, the C-H bond has a bond energy of 435 kJ/mol in methane and 410 kJ/mol in ethane.

isoelectronic

• Positive ions are smaller than the neutral atom because fewer electrons are repelling each other. • Negative ions are larger than the neutral atom because more electrons are repelling each other. • Among isoelectric atoms and ions, the one with more protons has a smaller atomic radius. Have same number of electrons but diff number of protons in nucleus More charge in nucleus will do better job of overcoming repulsive acts of electrons