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AS Chemistry A OCR Summary, Post 2015 Spec
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Summary notes of the official OCR AS Chemistry A OCR Student Book by Rob Ritchie and Dave Gent. Made by an A* student.
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Chemistry ASModule 2 Foundations in Chemistry
Atoms, Ions and Compounds
Atomic Structure and Isotopes
• Relative: Subatomic particles have tiny masses so use relative masses instead of grams.
• Electron: An electron has about 1/ 1800th mass of a proton, positioned in the shell.
• Protons: Number of protons identifies element- atomic number, the number on top in periodic
table but on bottom in the other notation.
• Mass Number: Aka nucleon number.
• Isotopes: Atoms of the same element with different number of neutrons and mass number.
• Isotope Properties: Same number of electrons means same chemical properties as chemical
reactions involve electrons. Isotopes different physical properties e.g. boiling point.
• Ions: Cations are positively charged ion. Anions are negatively charged ion.
• Polyatomic Ions: A charge on polyatomic ion means number of electrons more/ less overall, not
for each atom in compound. Have brackets for (SO4)2- polyatomic ions. Ignore sub- script e.g. 3 in
CO32-.
• Constructing Equation:
- Cross multiply charges on ions of compounds- product.
- Remember some elements exist as diatomic e.g. N2+- reactant.
- Balance equation normally.
- Add state symbols.
- See if need water, carbon dioxide or hydrogen to make up for missing elements in main
product.
Relative Masses
• Relative Mass: No units as it is ratio of two masses- it’s relative.
• Carbon-12 Isotope: One atom of carbon 12 is 12 atomic mass units. 1 atomic mass unit is 1/12th
of an atom of carbon- 12 = mass of a proton or neutron.
• Relative Isotopic Mass: It is the mass of an isotope compared to 1/12th of the mass of an atom of
carbon-12.
• Relative Atomic Mass: Ar. Is the weighted mean mass of an atom of an element compared to
1/12th of the mass of one atom of carbon-12. Shown in the periodic table. Most elements contain
mixture of isotopes.
• Calculating Ar: Sum of- percentage abundance of each isotope multiplied by their relative isotopic
masses. All divided by 100. Gives relative atomic mass.
,Formulae and Equations
• Ion Charges: Group 1 = 1+ ions etc. Group 7 = 1-.
• Need to know also Zn2+ and Ag+.
• Transition Metals: Form several ions with different charges. Ionic charge shown with Roman
numeral in name.
• Binary Compound: Compounds containing two elements. For the second element add suffix –ide.
Metal ion (+ve ion) always comes first.
• Polyatomic Ion: An ion containing more than one atom bonded together. Ammonium is NH4+.
Hydroxide is OH-. Nitrate is NO3-. Carbonate is CO32-. Sulphate is SO42-, Hydrogencarbonate HCO3-,
manganate (VI) MnO42−, phosphate PO43-
• Diatomic Molecules: Molecules which contain two atoms of same element bonded together. E.g.
hydrogen, nitrogen, oxygen, halogens, S8 and P4.
Amount of Substance
Moles
• Avogadro Constant: One mole = 6.02 x1023 particles, easier way to count number of particles. As
mole linked to carbon- 12. 1 mole in grams = relative atomic mass.
• Number of atoms = 6.02 x1023 x molar mass. 0.12 moles of S8 molecules means need to multiply
by 8 as actually working out molar mass, not moles.
• Moles (n) = Mass/ Molar Mass.
• Mass must be in grams. Molar mass in units g mol-1.
• How many molecules? Use Avogadro’s constant.
• How many bromine molecules present in each dm3 when bromine air is at a concentration of 4 x
10-6 gdm-3? 4 x 10-6/ 79.9 x 2 = mass/ Mr = 2.5 x10-8 moles. Then multiply by 6.02 x1023 to find
number of molecules.
• Remember bromine air or water is Br2 unless stated otherwise.
• A natural gas supply contains 16.0% H2S by volume. The H2S (g) in 1.50 × 108 dm3 of this natural
gas supply, measured at RTP, is processed into sulfur with an overall percentage yield of 95.0%.
Calculate the mass of sulfur, in g, obtained from 1.50 × 108 dm3 of natural gas supply.
- 0.16 x 1.50 × 108 = volume of H2S. Work out moles.
- Use molar ratio to find moles of S which is theoretical moles.
- Rearrange % yield formula to find actual moles of S. Find mass.
Determination of Formula
• Molecular Formula: Covalent compounds and elements exist as molecules.
• Empirical Formula: It is the simplest whole number ratio of atoms of each element in a compound.
Used for giant structures of metals and ionic.
• Relative Molecular Mass: Mr. Add relative atomic mass of each element making up molecule.
• Relative Formula Mass: Add relative atomic masses of elements in empirical formula.
• Calculating Empirical Formula: May need to do some subtraction/ addition if masses not needed
given.
- Find moles and simplify ratio by dividing by smallest number.
- Multiply if .5 or .7 (multiply by 3). If .2 can multiply by 5. Round up if can round up all of them.
, - Write empirical formula in order appeared in question.
• Calculating Molecular Formula:
- Find empirical formula and its Mr.
- Divide the Mr of the molecular formula by Mr of empirical formula.
- Multiply this factor by empirical formula.
• Knowing Compound but Not Element in the Compound:
- Divide Ar of element by Mr of substance to find percentage of element in whole compound.
- Multiply percentage by the mass of the compound to find out mass of the individual element.
- Use Mr Moles to find moles of element.
Hydrated Salts
• Water of Crystallisation: Water molecules that are bounded into crystalline structure of a
compound.
• Dot or .xH2O shows number of waters of crystallisation. Anhydrous means no waters of
crystallisation.
• Copper (II) Sulphate: When blue crystals of hydrated copper sulphate are heated, bonds holding
the water within the crystal broken. Crystalline structure is lost. Leaves white anhydrous copper
sulphate. Difficult to remove last traces of water, so can be very pale blue.
• Experiment: Weigh empty crucible. Weigh the crucible and the hydrated salt in it. Weigh crucible
and the anhydrous salt in it.
• When using hydrated salt in experiments, don’t know how much water is in the mass.
• Calculating Amount of H2O:
1- Find moles of anhydrous salt. Need to takeaway values.
2- Find moles of water. Need to takeaway values.
3- Find the simplest molar ratio e.g. 1: 5 by dividing moles of water by moles of salt. 5 would be
the amount of H2O molecules.
4- OR work out by subtracting molar mass of salt from salt + water compound and use remaining
molar mass to work out H2O number in front.
• Find Expected Mass of Anhydrous Salt: Find moles of hydrated salt. Use molar mass to find moles
of anhydrous salt.
• Converting Empirical to Dot Formula: Hydrogen is halved to be the number in front of H2O. Rest
of oxygen with other ion compound.
• Assumptions:
- All Water Lost: Some water may be left inside even if colours different. Difficult if both salt
forms, same colour. Must reheat salt repeatedly so mass remains constant which suggests all
water removed.
- Further Decomposition: Salts decompose further when heated. Copper (II) sulphate heated
very strongly, decomposes to black copper (II) oxide. Not as easy if no colour change.
Liquid- moles
• 1 cm3 = 1 ml. 1dm3 = 1000 cm3.
• Concentration: Unit mol dm-3. It is the amount of solute, in moles, dissolved in each 1 dm3 of
solution.
• Calculating Moles with Volumes: Moles (mol) = Concentration (mol dm-3) x Volume (dm3). If in
cm3/ 1000 = dm3.
,• Standard Solution: It is the solution of a known concentration. Prepared by dissolving an exact
mass of solute in a solvent and making up the solution to an exact volume.
• Calculating Standard Solutions: Steps can be opposite.
- Work out moles- using concentration volume of standard solution.
- Use Mr Moles to calculate mass of substance going in standard solution.
• Grams: Can have concentration with units g dm-3. Convert using Mr Moles formula. Value in
mol dm-3 x Mr = value in g dm-3.
Gas- moles
• Gas Volumes: At the same temperature and pressure, equal volumes of different gases contain
same number of molecules.
• Molar Gas Volume: Vm. It is the volume per mole of gas molecules. Volume of gas depends on
pressure and temperature.
• RTP: At RTP, 1 mole of gas molecules has a volume of 24.0 dm 3 = 24000 cm3. Therefore assume
the molar gas volume is 24.0 dm3 mol-1.
• Calculation- Moles: Moles = volume/ molar gas volume. Moles = Volume/ 24 dm3 mol-1. cm3/
1000 = dm3. Volume aka ‘volume of gas collected/ produced’.
• Ideal Gas Equation: Gas equation used when need to be more accurate or when carrying out
experiment not at RTP.
• pV = nRT.
- R always 8.314 J mol-1 K-1. n is moles- use moles of gas- or find using molar ratio.
- Volume is in m3. Cm3 to dm3 to m3, multiply or divide by 1000 each time.
- Pressure is in Pa. kPa to Pa is divide by 1000.
- C to K is +273.
• Assumptions for Ideal Gas: For an ideal gas, the assumptions for the molecules are
- Random motions
- Elastic collisions.
- Negligible size.
- No intermolecular forces.
• Experiment to Find Mr: Can use ideal gas equation to find Mr of volatile liquid. Add volatile liquid
to small syringe. Weigh small syringe. Inject sample into gas syringe. Reweigh small syringe to find
mass of volatile liquid added to gas syringe. Place gas syringe in boiling water bath. Liquid
vaporises to gas.
Reacting Quantities
• Stoichiometry: The ratio of moles of each substance in a chemical reaction.
• Balanced Equation: The balancing numbers gives molar ratio.
• Find Moles- Various States: Can have solid go to gas or liquid. If asks to calculate volume of gas,
remember there might be more than one gas.
- Work out moles of mass using Mr Moles.
- Molar ratio to find unknown moles.
- Concentration Volume for liquid or Volume/ 24 for gas.
• Identifying Unknown Metal: Attach gas syringe to conical flask. Add metal. Add HCl to flask and
quickly replace bung. Measure maximum volume of gas in the syringe. Use equations and molar
ratio to find unknown metal.
,• Theoretical Yield: It is the yield resulting from complete conversion of reactants into products.
Impossible as…
- Reaction may not have gone to completion.
- Other reactions may have taken place alongside main reaction.
- Purification of product may result in loss of some product.
• Percentage Yield:
- Find theoretical yield of product using values of reactant given and molar ratio.
- Find actual yield from values of product that are actually produced in question. Must use
appropriate mole formula to work these values out.
- (Actual y/ Theoretical y) x 100.
• Limiting Reagent: The reactant that is not in excess so will be used first and stop reaction. If don’t
know limiting reagent, need to find out by working out moles of each reactant and comparing with
equation. Calculations based on limiting reagent.
• Atom Economy: (Sum of molecular masses of desired products)/ (Sum of molecular masses of all
products) x100. Only worked out from balanced equation. Remember to include balancing
numbers when calculating molar masses.
• Only one product formed so 100% atom economy.
• Named waste product produced as well so atom economy less than 100%.
• Advantages- more readily available starting materials. No toxic products. High atom economy,
100% (no waste products). Compound is regenerated- used again.
• Sustainability: Sustainability depends on both atom economy and % yields.
• Low atom economy means poor sustainability so need to develop use for waste products to
increase atom economy. Therefore stop greenhouse gas emission as waste product.
• High % yield sustainable as efficient conversion from reactant to product. High atom economy
reduces amount of waste products, so increases sustainability.
Acids and Redox
Acids, Bases and Neutralisation
• Acids: Solid H2SO4 not an acid, when aqueous it is an acid. They ionise into (2)H+ and SO42. Need
to know acids- HCl, H2SO4, H3PO4, HNO3 and CH3COOH.
• Strong Acid: When dissolved in water, a strong acid completely dissociates in aqueous solution,
releasing H+.
• Weak Acid: It partially dissociates in aqueous solution. The equilibrium sign indicates forward
reaction complete.
• Bases: A compound acts as a base when it accepts H+. Base neutralises acid to form salt and water.
Bases are
- Metal oxides
- Metal hydroxides.
- Metal carbonates.
- Ammonia.
- NOT metals.
• Alkali: An alkali is a base that dissolves in water releasing OH- into solution. NaOH (aq) -> Na+ (aq)
+ OH- (aq). Need to know alkalis- NaOH, KOH and NH3.
,• Why Described as Salt: A salt is when the H+ in an acid is replaced by a metal ion or an ammonium
ion in neutralisation. Name the acid and metal ion. Chloride, sulphate, nitrate and ethanoate salts.
• Can divide moles of alkali reacted with moles of acid reacted, to work out number of H in chemical
formula or H+ ions replaced by metal ions to form salts.
• H4N2O3 made of NH4+ and NO3- ions
• Ionic Equation:
- Anything ionic/ aq, divide into ions. If balancing number for compound is 2, have the balancing
number for both ions as 2 as well.
- Include elements or covalent compounds as they are.
- Any ions that haven’t reacted are spectator ions and get rid of them.
• Acid and Alkali: Ionic equation for neutralisation reaction always H+ (aq) + OH- (aq) -> H2O (l).
• Acid and Base: Ionic equation is 2H+ (aq) + O2- (s) -> H2O (l). Remember to have all state symbols
as (aq).
• Acid and Metal Carbonate: H2SO4 (aq) + Na2CO3 (aq) -> Na2SO4 (s) + CO2 (g) + H2O (l). 2Na+ and
CO32- are the ions.
• Acid and Metal: 2HNO3 + Mg (s) -> Mg(NO3)2 (aq) + H2. Mg2+ and NO3- are the ions.
Acid- Base Titrations
• Titrations: Technique used to accurately measure the volume of one solution that reacts exactly
with another solution. Used to find concentration of solution, identify unknown chemicals and
find purity of substances.
• Volumetric Flask: They make up standard solutions very accurately due to small neck.
• Preparing Standard Solution:
1- Solid weighed. Solid dissolved in beaker using some distilled water.
2- This solution transferred to volumetric flask. Last traces of solution rinsed into flask with
distilled water.
3- Flask filled to graduation line by adding distilled water a dropwise until bottom of meniscus
lines up at eye level. If too much water added, solution becomes dilute and must be prepared
again.
4- Volumetric flask inverted to mix solution. Or results would be inconsistent.
• Burette: Rinse burette with solution- pour out by slow rotating and tilting. When filling burette,
run excess solution through tap to remove air bubbles.
• Acid- Base Titrations:
1- Add one solution using pipette and indicator to conical flask.
2- Add other solution to burette and record initial burette reading to nearest 0.05cm3.
3- Run solution in burette into solution in conical flask, swirling conical flask.
4- Eventually indicator changes colour at end point of titration. Record final burette reading. The
volume of solution added from the burette called titre (subtract initial and final burette
reading).
5- Trial titration to find approximate titre (cm3).
6- Titration then repeated, adding solution dropwise as the end point approached. Repeat
titration until two titrations are concordant within 0.1 cm3.
7- Calculate mean titre from concordant titres. Reject ones not agreeing within 0.1.
• Identification of Carbonate: X2CO3.
- Take measurements of mass of weighting bottle to find mass of X2CO3 solid powder. Prepare
250cm3 solution of unknown carbonate in volumetric flask.
, - Using pipette measure 25cm3 prepared solution into conical flask. Using burette, titrate this
solution using HCl. Work out mean titre.
- Calculate amount of HCl in moles reacted using CV where V is mean titre used.
- Use molar ratio to work out amount of unknown carbonate reacted in moles.
- Scale up moles by multiplying this mole value to get original solution prepared- divide to get
1cm3 and then x 250 to get 250cm3.
- Find molar mass in gmol-1 using moles of 250cm3 solution and mass of X2CO3 using Mr moles.
- Use molar mass and take away molar mass of CO3, to leave X2 and use period table to find
unknown element.
Redox
• Oxidation Number: Aka oxidation state. Oxidation number sign placed before the number, unlike
ionic.
• Rules in Most Cases:
- Elements- Oxidation number is zero. Even Cl2.
- Most Metals- Oxidation number of group 1, 2 and 3 same as ionic charge, +1 or +2 or +3. Look
at periodic table, to check none of elements are in the metal groups.
• Special Cases:
- Metal-
- H in metal hydride (NaH) is -1.
- Al, a group 3, is always +3.
- F is always -1.
- Non- Metal-
- Oxygen usually -2.
- O in peroxides (H2O2) is -1 as H is +1.
- O bonded to F is +2.
- Cl, Br and Iodine usually -1.
- Cl, Br and I +1 when bonded with F or O.
• Compounds: Sum of oxidation numbers = total charge of compound. Ones above are always that
number, so can work out unknown. Metals take priority over non-metals when assigning.
• The oxidation number represents one atom of an element. The total change will depend on how
many atoms of the element in a compound there are- sub script.
• Roman Numerals: Roman numeral shows the oxidation number of element. As some elements
have more than one ion charges.
• OILRIG: Oxidation is the loss of electrons and increase in oxidation number. Reduction is gain of
electrons and decrease in oxidation number. Both happen in one equation so one element will be
reduced and another will be oxidised. This is called a redox reaction.
• If oxidation number has increased by 2 and has been oxidised, means has lost two electrons.
• What has been oxidised and reduced: Oxygen has been oxidised as oxidation number has
increased from 0 to +2. Fluorine has been reduced as oxidation number has decreased from 0 to
-1. Total changes in oxidation numbers balance so include sub script.
• Half Equations: Iron loses electrons 2Fe -> 2Fe3+ + 6e-. Chlorine gains electrons 3Cl2 + 6e- -> 6Cl-
The electrons gained and lost balance.
• Redox Reactions of Acids: Metal + acid -> salt + hydrogen.
Zn (s) + 2HCl (aq) -> ZnCl2 (aq) + H2 (g).
2Al (s) + 3H2SO4 (aq) -> Al2(SO4)3 (aq) + 6H2 (g).
,• Named metal dissolves. Bubbles.
Electrons and Bonding
Electron Structure
• Shells are energy levels.
• Atomic Orbitals: It is region around a nucleus that can hold up to two electrons, with opposite
spin. Must have opposite spin due to repulsion between two negative electrons.
• Four Different Atomic Orbitals: s-, p-, d- and f-. The greater the shell number n, the greater the
radius of its orbital. Must write in lower case and write superscript accurately.
• S- Orbitals: The electron cloud is within the shape of a sphere.
• P- Orbitals: Electron cloud is in shape of dumb- bell. There are three p- orbitals at right angles to
one another, called px, py and pz.
• d- Orbitals and f- Orbitals: From n = 3 there are five d- orbitals. From n = 4 there are seven f-
orbitals.
• Sub- Shell: Within a shell, orbitals of the same type are grouped together as sub- shells. The
number of orbitals increases with each new type of orbital so number of electrons in sub shell
increases.
• A 3p orbital would be 2, not 6. The 3rd shell would be 18. The 3s sub shell would be 2.
• Filling Orbitals: Orbitals fill in order of increasing energy. 3d sub- shell is at a higher energy level
than 4s sub- shell so 4s sub- shell fills first like 3p, 4s, 3d.
• Electrons in Box Modal: Two arrows represents two electron, indicating opposite spin down. One
box represents one orbital.
• Energy of Orbital: Energy is equal in orbitals. One electron occupies each
orbital before pairing starts to prevent further repulsion between paired
electrons. In this example there are one completely filled orbitals.
• Watch out for ionic charges on elements, don’t just see the element and go
straight in.
• Electron Configuration: List number of electrons occupying orbital in superscript e.g. 1s2, 2s2, 2p6…
3d10, 4p6. 3d10 is before 4s2 as electron configuration shows shell order NOT order of filling or how
it appears on periodic table. All numbers are together.
• Shorthand Electron Configuration: Use previous noble gas in periodic table e.g. for Krypton [Ar]
4s1.
• Periodic Table: Periodic table can be divided into blocks corresponding to highest energy sub-
shell. All the s- and p- blocks correspond with period number. The d- block are p minus 1.
,• Electron Configuration of Ions: When forming ions, the highest energy sub- shells loses or gains
electrons first. Once 4s is filled, 3d energy level falls below 4s energy level so 4s sub- shell empties
before 3d sub-shell as in order of electron configuration.
Ionic Bonding and Structure
• Bonding Questions: E.g. melting point question ‘in terms of bonding and structure.’
- Name structure.
- Name bond.
- Write what the attraction is between e.g. + metal ions and –ve delocalised electrons in
metallic, atoms in covalent bond and molecules in simple molecular.
- Also can define if not done already.
• Ionic Bonding: Is the electrostatic attraction between oppositely charged ions.
• When given table and comparing melting points, don’t go too in depth. If numbers are similar,
that’s good enough. If melting point much larger, then state if giant covalent or giant ionic.
• Ions: Common cations are metal ions and ammonium ions. Common anions are non- metal ions
and polyatomic ions. Ions formed have outer shells with same electron configuration as nearest
noble gas.
• Dot and Cross Diagram: Only outer shells shown but for full outer shell leave empty and just draw
a circle. In ionic bonding draw more than one ion if needed, e.g. two magnesium ion for every
oxygen. Don’t draw squiggly line like I did at GCSE. But remember ionic charges.
• Don’t assume covalent when see dot and cross, check periodic table.
• Structure of Giant Ionic Lattice: Each ion is surrounded by oppositely charged ions in the
structure. It is a regular cubic arrangement. Don’t write ionic, write giant ionic.
• High Melting and Boiling Point:
- Ionic compounds need high temperatures to provide larger quantity of energy.
- To overcome strong electrostatic forces of attraction between oppositely charged ions.
• Greater Ionic Charge and Larger Size:
- Lattices have higher melting point
- due to stronger attraction between ions when…
- Smaller ion size.
- Greater ionic charges.
• Soluble: Many ionic compounds dissolve in polar solvents such as water. The ionic lattice breaks
down as water molecules surround the positive and negative ions.
- The positive ion is attracted towards the oxygen of the water molecule as it is δ-.
- The negative ion is attracted towards the hydrogen of the water molecule as it is δ+.
- Lattices with ions with greater ionic charges will be too strong for water to break it down so it
will not be soluble.
• Electrical Conductivity as Solid: Ions are fixed in position in an ionic lattice so are not mobile in a
solid. Cannot conduct electricity.
, • Electrical Conductivity as Liquid, Molten or Aqueous: Solid ionic lattice breaks down. Ions can
move and conduct electricity when molten, have mobile ions.
NH4Cl an ionic compound
Covalent Bonding
• Covalent Bonding: It is the strong electrostatic force of attraction between a shared pair of
electrons and the nuclei of the bonded atoms. Covalent bonding occurs between non- metallic
elements and polyatomic ions.
• Orbital Overlap: A covalent bond is the overlap of atomic orbitals. The bonded atoms often have
same electronic structure as nearest noble gas.
• Attraction: Attraction is localised as attraction acts solely between shared pair of electrons and
nuclei of bonded atoms unlike ionic.
• Lone Pairs: Show lone pairs when drawing- an outer shell pair of electrons that is not involved in
chemical bonding.
• Carbon forms 4 bonds. Nitrogen forms 3. Oxygen forms 2. Hydrogen forms 1.
• Boron: It has three outer- shell electrons. Forms covalent compounds like BF3 which has six
electrons in the outer shell altogether and not eight.
• Phosphorus, Sulphur and Chlorine: Phosphorus has five outer- shell electrons and can form PF3
and PF5. Sulphur can form SF2, SF4 and SF6. Chlorine can form ClF, ClF3, ClF5 and ClF7.
• Expansion of the Octet: All three above are in Period 3, n=3, so can hold maximum of 18 electrons.
It called expansion of the octet and only possible from the n=3 shell- the d sub shell.
• Can arrange their electrons differently by having some electrons as single and some as lone pairs.
• Dative Covalent: Aka coordinate bond. It is a covalent bond in which both pair of electrons has
been donated by one of the atoms. The shared pair of electrons was originally a lone pair of
electrons on one of the bonded atoms. A dative covalent bond shown by a → in displayed formula.
• Use normal stuff first and if that doesn’t work, use Dative and n=3 holds 18e- as last resort to make
stable compound.
• Multiple Covalent Bond: In double covalent, electrostatic force of attraction between two shared
pairs of electrons and the nuclei of bonded atoms.
• If stuck drawing compound, start drawing dot and cross diagrams from different element.
• Average Bond Enthalpy: The larger the value of this, the stronger the covalent bond.
Shapes and Intermolecular Forces
Shapes of Molecules and Ions
• Electron- Pair Repulsion Theory: Electron pairs arranged as far apart as possible to minimise
repulsion. Number of bond pairs and lone pairs of electrons surrounding central atom determine
shape. This holds the bonded atoms in a definite shape.
- Name shape.
- State number of bonded pairs of electrons and lone pairs of electrons.
- If a lone pair state electron pairs repel but lone pair repels more than bonded pairs.
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