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Honors Chem 1st and Second semester (CHEMICAL EQUATIONS (Double…
Honors Chem 1st and Second semester
PROPERTIES OF MATTER
PHYSICAL & CHEMICAL CHANGES
Physical Changes
are changes that affect the physical properties of the matter of a substance. For example = Cutting a piece of paper
Chemical Changes
are changes that affetc the physical state of matter of a substance, meaning a new substance is made. For example = rusting iron, rust is created as a byproduct of the chemical change that occured
Physical and Chemical Properties
Physical properties
can be determined without changing the
nature of a substance. for example = the color of a substance would be a physical property
Chemical Properties
are only identified by the cause of a chemical reaction. For example = a substance's ability to rust would be a chemical property, Ability to react with oxygen etc.
ATOMIC STRUCTURE
Atomic Theory's
Daltons Atromic Theory
=
Daltons Atomic theory stated that all matter was composed of extremely tiny particles called atoms. All the atoms of an element will hold the same properties, such as same size, mass etc and atoms of different elements will have different properties.It stated that atoms cannot be created destroyed or
Subdivided
. Atoms of different elements combine in simple whole number ratios to form compounds.
Modern Atomic Theory
=
The modern Atomic Theory is very similar to Daltons Atomic Theory in the way that all matter is composed of matter, atomns of an element differ in properties to other atoms i of other elements. However, Atoms cannot be subdivided, created or destroyed in ORDINARY REACTIONS, but these changes can occur in nuclear reactions.. Atoms of an elements have a charecteristic average mass that is unique to the element.
Structure =The atom is composed of three types of subatomic particles
Neutrons= These subatomic particles are located in the nucleus of an atom. It has no charge and these particles make the atoms of each element have different mass( different number of neutrons). These atoms with same number of protons but different number of neutrons are called
isotopes
. The atomic mass ( sum of neutrons and protons) of the element will be the average of the masses of these isotopes.
Electrons = Located in orbitals surrounding the nucleus of an atom, it has a negative charge and in neutral elements the number of protons= number of electrons :
Discovered by JJ Thompson, in 1897. he used a cathode ray(they pass electricity through low pressure contained gas) to deterimine the position of the electron in the atom. FDrom this study, JJ Thomson decuced that all atoms have identically charged electrons and that atoms are neutral because the postive charge(protons) and the negative charges(electrosn) cancel each other out. Electrons have little mass.
Protons
= These subatomic particles are located in the nucleus and have a positive charge. they make up the properties of the element. The atoms ofa given element will always have the same number of protons. The Atomic number ( denoted as Z) of an element is the number of protons it has
Discovered by Rutrherford's Gold Foil experiment, in whcih alpha particles ( helium nuclei) were fired at a thin sheet of gold. Many particles passed through, few deflected while very few were
greatly
reflected, proving that the nucleus is small, dense and positively charged
Two types of atomic structure models
Quantum Model
in this model, the elctrons do not have definite orbits and are pure charges floating somewhere in the orbital area.
Bohr model
, called the planet model, because it looks like the elctrons in orbits around the nucleus are planets circling the sun( nucleus)
Nuclear Chemistry = The study of the structure of atomic nuclei and the changes it undergoes.
Radioactivity
= Three types of radioactive particles
Beta Radiation
= beta particles that are similar to electrons with a -1 and a mass of 1/1837 amu ( usually regarded as 0). Relatively strong radiation that can be stopped by a thin sheet of metal. Symbol =e
Gamma Radiation
= Pure,high energy electromagnetic radiation that has no mass and no charge. It is the strongest form of radiation and so can only be stopped by lead or concrete. Symbol = y
Alpha Radiation
= Alpha particles are helium nuclei, having 2 protons and 2 neutrons. It had a 2+ charge and mass of 4 amu. Lightest radiation, it can be stopped by a sheet of paper or skin. Symbol = H, or a
NUCLEAR REACTIONS
= Isotopes of one element are changed to the isostopes of another element. The content of the nucleus changes and large amounts of energy releases during these reactions
Alpha Beta and Gamma Decay
Alpha Decay
occurs when an unstable nucleus emits a particle with 2 protons and 2 neutrons. This radioactivity is the most ionizing but least penetrating. In this type of decay the atomic number decreases by 2 amu and atomic mass decreases by 4 amu
Beta Radiation
occurs when an unstable nucleus emits an electron and this changes the nucleus (the protons turn into neutrons).This radiation has half of alpha radiations ionizing ability but is 10 times more penetrating, Atomic number increases by 1 amu while atomic mass remains the same.
Gamma Radiation
=High energy photons of light, no loss of particles form nucleus and there is no effect on the nucleus. Least ionizing radiation but most penetrative. This radiation usually occurss after the nucleus undergoes some other type of decay.
Half Life
= The time required for a radioactive substance to decay half of its original amount.
The formula is Mass remaining = Initial mass *(0.5)^ (time/half life).
To solve for time or half life the formula is , time/half life =[ Ln(Final Mass) - Ln(initial mass) ] /Ln(0.5)
Bonding
, Occurs between cations and anions
Why do atoms bond?
Chemical Bonds form because they lower the potential energy between the charged particles that comprise electrons. Bonds dorm when the potential energy of bonded atoms is less than the potential energy of the individual atoms. Bond formation results in a loss of energy while breaking a bond takes a gain of energy,
Nucleus to electron repulsion
Electron to Electron repulsion
Nucleus to nucleus repulsion
Three types of bonds
Ionic Bonds
= form between metals and non metals . The electrons are transferred. for example = NaCl , Sodium has to give 1 electron away while Chlorine need to take 1 electron so they exchange electrons and form a compound
Metallic bonds
= These form between two metals, the two metals share their electrons in a pool of electrons called the ' sea of electrons' . This explains why metals are so conductive ( fluidity of electrons)
Covalent bonds
= from between non metals only, electrons are shared between each other. This bonding followes octet and HONC rule
Non polar covalent
=when electrons are shared equally. Electronegativity differenceis btween 0 to 0.3. eg =CH4
Polar covalent
= when electrons are unequally shared, the difference is between 0.3 and 1.7
Hybridization
is combining two or more orbitals or nearly equal energy within the same atominto orbitals of equal energy. Steric number
2
= sp,
3
= sp^2 ,
4
= sp^3
VSEPR and Molecular Gemetery
Lewis Structures
= Has to follow the octet rule
HONC
rule,
H
stands for hydrogen and halogens only forming 1 covalent bond,
O
stands for oxygen and sulfur having the ability to form to 2 covalent bonds.
N
for nitrogen and phosphorous having the ability to form 3 covalent bonds and
C
stands for Carbon and Sillicon having the ability to form 4 covalent compounds
Steps to draw Lewis structures = first count total number of valence electrons and divide it by 2. Put the least electronegative atom as the central atom and place all the other atoms around it. Place a pair of electrons between central atoms and all other atoms and distribute the rest.. Remaining pairs must be put next to the central atom. Resonance occurs when there is more than 1 valid lewis structure for a molecule
VSEPR
stands for
V
alence
S
hell
P
air
E
lectron
R
epulsions. The structure is determined by the minimizing electron pair replusions. This is done by finding the steric number using the AXE method
AXE
method =
A
represents the cemtral atom.
X
represents the number of bonds formed between the central atom and surrounding atoms.
E
represents the number of lone pairs present on a central atom. the STERIC NUMBER is the sum of
X & E
Molecular shapes= the main shapes are linear, trigonal planar, bent or angular, tetrahedral, pyramidal(many types), square planar, t shape, seesaw.
Periodic Table
Periodic Trends
Ionization energy
= energy required to remove an electron from an atom in a gaseous state. It tends to increase accross a period ( more electrons, the more attraction and therefore more energy needed to remove an electron. It tends to decrease down a group (more energy levels)
Electronegativity
= measure of the ability of a compound to attract electrons.. It tends to increase accross a period( as radius decreases, electronegativity increases) and decreases down a group( as radius increases, electronegatvitiy decreases)
Atomic Radius
= 1/2 the distance between nuclei in a covalently bonded diatomic meolecule.. Tends to decrease accross a period( Zeff increases, atomic radius decreases) and it tends to increase down a group (lthe less electronegativity, the more atomic radius
Groups
, the periodic table is mostly grouped into metal, transition metals, metalloids, non metals, halogens and noble gasses
Electron Configuration and Quantum Numbers
Electron Configuration
Electron Orbitals = An orbital is a region in an energy level where there is a high probability of an electron being found. The shapes of the orbitals are defined as the surface that holds 90 percent fo total electron probability.There are 4 types.
p
orbital shape , there are 3 dumbell/double lobed p orbitals in each energy level above n = 1, each assigned to its own axis(x,y,z) in space. The p block is columns boron to neon, 6 colums represent it
s
shape has a sperical shape centered around the nucleus and the origin of the 3 axes in space., They fill the first two columns of the periodic table. 2 colums represent it.
d
orbitals have 5 orbitals that look like double dumbells, whole transition block represents it (10 columns)
f
orbitals have 7 orbitals and 14 columns reprenting it.
The purple area are the s orbitals, the yellow areas are the d orbitals, the green areas are the p orbitals and the blue areas are the f orbitals
Electronic Configuration RULES
Pauli Exclusion Principle
states that no two electrons can have the same quantum number. It also states that only two electrons can fill a single orbital and they must have opposite spins
Hunds Rule
states that electrons must fill each equal energy orbital first before pairing up with electrons in single orbits.
Aufbau Principle
, This priciple states that electrons must fill the lowest energy levels first before filling higher energy levels.
For example , He = 1s^2
Quantum Numbers,
Quantum Mechanical model of the Atom
-= Mathematical laws that identify the regions where electrons are most likely to be found.
Heisenberg Uncertainty Principle
= States that the postion and the speed of an electron cannot be determined simultaneously
Each electron has 4 specific quantum numbers
Angular Momentum number
l
denotes orbital in which the electron is located. s orbital = 0, p orbital =1, d orbital = 2, f orbital = 3
Magnetic Quantum number, denoted by
m
is the orientation of electron orbital with respect to 3 axes in space
Principal Quantum Number
n
, this denotes the location of the electron
Spin quantum number
s
denotes the direction of spin of an electron with a magnetic field. 1st half is +1/2, 2nd half is -1/2
For example phosphorous [3,+1,+1,+1/2]
Mixtures
Homogenous = uniform in composition. eg = salt water
Heterogenous = not uniform in composition. eg = sand + water
CHEMICAL EQUATIONS
Double Replacement
When the ions of 2 compounds exchange places in an aqueous solution forming 2 new compounds. Eg = AX +BY = AY + BX
Combustion
= When a substance reacts with oxygen to always form carbon dioxide and water and releasing a large amount of energy and light
Single Replacement
= When a higher reactive substance replaces another element in a compound. AX + B = A + BX
Decomposition
= When a compound breaks into smaller simpler substances. For example = AX = A + B
Stoichiometery
is determining the desired quantitative data of a substance using the relationship between the reactants and the products in a chemical reaction
Based on
Law Of Conservation Of Mass
that states that matter cannot be created or destroyed in a chemical reaction. A balanced equation helps with this
Mole/Mass
The coefficients of the balanced equation establish the mole ratio. Mole to gram conversion , Mass of substance/ moles of substance = Molar mass of substance / 1 mole
States of Matter
Kinetic Molecular Theory
1)Gases consist of tiny particles far apart relative to its size.
2)Collisions between particles and particles, and particles and the walls of the container are elastic collisions.
3)No kinetic energy is lost in elastic collisions.
4) Gas particles are always in constant rapid motion
5) There are no forces of attraction between gas particles
Ideal gases fit these requirements. However this is theoretical, most gasses are in fact actually real gasses. Ideal gasses are more likely to be small non polar molecules with
high
temperatures and
low
pressure while real gasses are more likely to be polar molecules with
low
temperature and
high
pressure.
Gases
Properties
Gases is the state of matter that has the least density, least molecules in a space. They have the lowest density of all states of matter, expand to fit the walls of the container. They diffuse and effuse very quickly. They are very fluid.
Pressure is Force per unit area. Its unit is Pascal.
Gas Laws
Gay Lussac's Law : P1/T1 = P2/T2
Combined Gas Law: P1V1/T1 = P2V2/T2
Charles law: V1/T1= V2/T2
Ideal Gas Law: PV = nRT
Boyle's law : P1V1= P2V2
Daltons Law of Partial pressure P(total) = P1 + P2 + P3 .......
Avogadros Law : V1/ n1 = V2 /n2
Phase Changes
Solid to liquid: Melting ( energy absorbed)
Liquid to gas: Vaporization( energy absorbed)
Gas to liquid: Condensation (energy is released)
Liquid to solid: Freezing ( energy absorbed
Gas to solid: Sublimation ( energy is released)
Solid to gas: Deposition ( energy absorbed
Solutions
Types of solutions=
Saturated+ solution with maximum solute dissolved in it
Super Saturated : Solution with more than maximum solute dissolved.
Unsaturated : Solution with no solute dissolved in the soltion.
Energy and Chemical Changes
Types of energy
Joule: SI unit of energy.
Calorie :4.184 J is 1 calorie. It is the heat required to raise the temperature of 1 gram of water to 1 Celsius. Specific heat is the heat required to raise 1 gram of any other substance than water
Latent heat of Phase Change
Heat to vaporization : Energy that must be absorbed/ released in order to convert 1 mole of liquid to gas or vice versa at its boiling point. unit = 40.7 kJ/ mol or 2240 J/g .
Heat of fusion : Energy that must be absorbed/released in order to convert 1 mole of solid to liquid or vice versa at its melting point. unit = 6.01 kJ/mol, 333 J/g.
Universe = system + surroundings
Reaction rates - Kinetics
Collision theory
This theory states that a reaction can only occur if the particles hit each other with the correct amount of energy and proper orientation
Endothermic Reaction
Exothermic reaction
Factors affecting collision theory
Increased temperature
Increased Surface area
Catalysts: These speed up the reaction, they provide a shortcut of sorts for the raction
Increased pressure
Rate laws ; Law o R = k(a)(b)
Chemical equilibrium
equilibrium is the balance of forces
Le Chatelier's Principle : This principle states that when outside factors affect an equilibrium like temperature or pressure, the reaction will shift accordingly .
The equilibrium constant is a ratio of the concentration of the products to the concentration of the reactants
Acids and Bases
Acids: have a pH of lower than 7. They donate hydrogen ions. They are sour
Strong acids: They dissociate completely with/ without water to form conjugate base and a proton
weak acids : dissosiate in equilibrium. , they form in equilibrium. small k = reactant favoured. the equilibrium constant is Ka ( follows law of mass action). pKa = -log(Ka)
Bases: They have a pH of higher than 7. They donate hydroxide ions.
Strong bases: completely dissociate in water
Weak bases:dissociate in equilibrium. basic constant is Kb. Kb
Ka = Kw. Kw is the constant for water, 1.0
10 to the power 14.
5 percent rule
1 liter of soltion = 55 moles of water
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