Chapter One

  • Matter is anything that takes up space.
  • Solids and liquids are incompressible.
  • Gases are highly compressible.
  • Gas refers to a substance that is a gas at room temperature; vapor refers to the gas form of a solid/liquid substance.

Chapter Three

  • Physical properties are characteristics that can be observed without changing composition.
  • Extensive properties depend on the amount of a substance while intensive does not.
  • Density is an example of an intensive property.
  • A chemical property is the ability of a substance to combine with or change into other substances.
  • A phase change is the transition of matter from one state to another.
  • Five signs of a chemical reaction include:
    1. Colour Change
    2. Formation of Gas (Bubbles)
    3. Formation of Precipitate (Solid)
    4. Evolution of Heat/Light
    5. Formation of Water
  • The Law of Conservation of Mass states that the mass of the reactants is equal to the mass of the products.
  • Substances are mixtures or pure-substances.
  • Elements/water are pure because they cannot be physically separated.

Mixtures:

  • In a homogenous mixture/solution, individual items cannot be seen; in a heterogeneous mixture, the difference can be seen.
  • Colloids are heterogeneous.
  • A solution is made up of a solute and a solvent. Water is always a solvent.
  • Solutes are what is dissolved; the solvent is what the solute is dissolved into.
  • Alloys are mixtures of metals.
  • Filtration, distillation, crystallization, sublimation, and chromatography are methods that can separate mixtures.
  • A solvent is saturated when it cannot dissolve anymore solute.
  • A compound is comprised of two different elements bonded together.
  • A molecule is a covalent bond between non-metals.
  • The Law of Definite Proportions states that a compound is always composed of the same proportions of elements no matter the sample size.
  • Percent by Mass (%)=(mass of element)/(mass of compound) x100

Naming Compounds

Covalent Naming:

  • All elements in a covalent bond are non-metals.
  • There is a prefix system where each element is given a prefix based on the number of those atoms in the chemical formula. It follows:
  1. Mono
  2. Di
  3. Tri
  4. Tetra
  5. Penta
  6. Hexa
  7. Hepta
  8. Octa
  9. Nona
  10. Deca
  • ___Element One ___Element Two-ide
  • In front of oxygen, drop a’s and o’s.
  • Never use mono for the first element.
  • Br I N Cl H O F always occurs in sets of two.

Ionic Naming:

  • Ionic bonds occur between a metal and a non-metal.
  • Write the name of the first element then add ide to the second element.
  • Ionic never uses prefixes!

Polyatomic Ions

  • These are chemical bonds between two elements. Unlike compounds, all polyatomic ions have charges.
  • Prefixes of per and hypo can be used on polyatomic ions; suffixes of ate and ite can also be used.

Perchlorate=ClO4

Chlorate=ClO3

Chlorite=ClO2

Hypochlorite=ClO

Measuring/Sig Figs

  • When measuring, always measure one place past the smallest marking on the device.
  • Read graduated cylinders at the bottom of the meniscus.
  • The answer to a problem cannot be written with more or less significance than the numbers used.
  • Always use the given number with the least amount of significance.

Significant Figures Rules:

  1. All non-zero numbers are significant.
  2. All zeros between significant numbers are significant. (Sandwich Rule)
  3. All zeros that are to the right of the decimal and to the right of a significant figure are significant.
  4. All zeros with a line over them are significant.
  5. All other zeros are not

Atomic Theory

  • The nucleus is very small and contains the atom’s protons and neutrons. They are packed very tightly together and make 99.99% of an atom’s mass.
  • The electron cloud takes up most of the volume but little mass. Electrons are very spread apart.
  • Aristotle believed that matter was continuously divisible.
  • In 1794, Proust published the Law of Definite Proportions.

Dalton’s Theory:

  1. All matter is composed of atoms.
  2. All atoms of a given element are identical in size, mass, and other properties.
  3. Atoms cannot be subdivided, created, or destroyed.
  4. Atoms of different elements combine in simple whole number ratios to form compounds.
  5. In chemical reactions, atoms are combined, separated, or rearranged but not created or destroyed.
  • Electrons were first discovered in 1897.
  • In 1909, Millikan calculated the charge and mass of electrons using his Oil Drop Experiment.

Cathode Ray Experiment:

  • Thomson and Crookes used a vacuum tube that contained a low pressure gas.
  • At each end of the tube, they placed positive and negative electrodes that were connected through the glass to the outside of the tube.
  • Led to the discovery of the electron.
  • Thomson and Crookes hooked the electrodes up to a power source and observed that rays traveled from the negative electrode (cathode) to the positive electrode (anode).
  • Because the rays originated at the cathode, they called them cathode rays.
  • To find out more, they performed some simple experiments.
  • First, they exposed the rays to a positive plate and observed that the rays bent towards the plate.
  • Negative plates bent the rays away.
  • They charged the gasses and metals and got the same results. Proposed that particles were in everything.

Models:

  • Thomson came up with the Plum Pudding Model.
  • Millikan showed that electrons have a mass and a negative charge.
  • Rutherford attempted to prove that the plum pudding model was correct; he proved it wrong.
  • Rutherford conducted the Gold Foil Experiment. He noticed that two percent of alpha particles were deflected.
  • Rutherford came up with the Planetary Model and the nucleus.
  • Mass listed on the Periodic Table is the average mass of all of the element’s naturally occurring isotopes.
  • Radiation occurs when the nucleus is falling apart. It emits particles or energy.

Quantum Theory

  • Ground state is the lowest possible energy state an electron can occupy.
  • Hydrogen gives off the Lyman, Balmer, and Paschen series.
    • The Balmer series is visible light given off by electrons jumping to energy level two. This gives off four frequencies.
    • The Lyman series is ultraviolet light given off by electrons going to level one.
    • The Paschen series is infrared light given off by electrons jumping to level three.
  • De Broglie described electrons as waves.
  • Heisenberg said that it is impossible to know the velocity and position of an electron at the same time. (Heisenberg Uncertainty Principle)
  • Schrödinger came up with the Quantum Mechanical Model.

Orbitals:

  • Electrons have s, p, d, and f energy sub-levels.
  • S are sphere orbitals.
  • P are dumb shell shapes along the x, y, and z axises.
  • D has its first four orbitals as cloverleaves along different axises. The other is uniquely shaped.
  • F are seven uniquely shaped orbitals.
  • Up to two electrons can fit in any level.

Electron Configuration:

  • Electron configuration is where electrons probably are in an atom.
  • The Aufbau Principle states that electrons occupy the lowest energy orbital available.
  • The Pauli Exclusion Principle states that only two electrons can occupy an orbital.
  • Hund’s Rule states that due to repulsion, unoccupied orbitals are inhabited first.
  • The Half Full/Full Rule states that it an element is one electron away from being either full or half full in d or f. an electron is pulled down into the d/f orbital.
  • When writing orbital notation, up arrows are drawn first.
  • The terminal electron is the last electron drawn in.

Quantum Numbers:

  1. The Principal (n) tells you the energy level of an electron.
  2. Angular Momentum (Azimuthal) is the sub-level of an electron. (s=0, p=1, d=2, and f=3)
  3. Magnetic (ml) tells you the orientation of the orbital (which orbital). The middle is zero with negatives to the left and positives to the right.
  4. The Spin (ms) tells the spin of an electron. (↑=+½, ↓=-½)

Chapter Six

  • Lavoisier compiled a list of known elements and placed them together based on properties.
  • Newlands organized elements by mass and came up with the Law of Octaves.
  • Meyer and Mendeleev noticed connections between mass and properties. Mendeleev is the “father” of the Periodic Table.
  • Moseley organized the table by proton number.
  • Groups are vertical; periods are horizontal.

Groups:

  • B, Si, Ge, As, Sb, Te, Po, and At are metalloids.
  • Alkali metals (Group One) are the most reactive elements and do not exist in free form in the Earth’s crust.
  • Alkaline Earth metals (Group Two) are not found in free form in the Earth’s crust.
  • Group Sixteen elements are known as “Calcogins”.
  • Group Seventeen elements are known as “Halogens”. They are the most reactive non-metal group.
  • Group Eighteen elements are known as the noble gases and are very unreactive.
  • Metals are cations; they lose
  • Non-metals are anions; they gain
  • Fr is the most reactive metal; F is the most reactive non-metal.

Periodic Table Trends:

  1. Energy Level– Higher energy level is further away.
  2. Effective Nuclear Charge– More charge pulls electrons in closer.
  3. Electron-Electron Repulsion (Shielding Effect)
  • If you are not a valence electron, you are a core electron.
  • Energy levels and shielding have an effect on group trends; nuclear charge has an effect on periodic trends.
  1. Atomic Size– Increasing atomic number causes atoms to gain energy levels and increases in size. Going left to right (across a period) causes atoms to get smaller due to nuclear charge pulling valence electrons closer.
  2. Ionization Energy– Amount of energy required to remove an electron from a gaseous atom. Energy required to remove only the first electron is called first ionization energy. Second ionization energy is greater than the first and so on… Ionization energy is determined by nuclear charge and distance. A greater charge means greater ionization energy; greater distance means less ionization energy. Filled and half-filled orbitals have a higher IE while one more than full or half-full has lower IE. As you go down a group, the first IE decreases because the electron is further from the nucleus and there is more shielding. All elements in the same period have the same energy level. They have the same shielding but have an increasing nuclear charge. IE generally increases from left to right, with exceptions at full and half-full orbitals.

Cations:

  • Cations form by losing electrons.
  • Isoelectronic species are atoms with the same number of electrons.
  • Cations are smaller than the atom they came from- not only do they lose electrons; they lose an entire energy level.
  • Cations of representative elements have the noble gas configuration before them.

Anions:

  • Anions form by gaining electrons.
  • Anions are bigger than the atom they came from-they have the same energy level but greater nuclear charge.
  • Anions of representative elements have the noble gas configuration after them.
  1. Ionic Size– Each step down a group is adding an energy level. Ions therefore get bigger as you go down because of the additional energy level. Across the period from left to right, the nuclear charge increases- so they get smaller. Notice the energy level changes between anions and cations.
  2. Electronegativity– Tendency for an atom to attract electrons to itself when it is chemically combined with another element. An element with a big electronegativity means it pulls the electron towards itself strongly. The further down a group, the farther an electron is away from the nucleus, and the more electrons an atom has makes the atom more willing to share (low electronegativity). For the period trend, metals are at the left of the table and easily let electrons go, thus low electronegativity. At the right are non-metals, they want more electrons. They try and take them away (high electronegativity).
  3. Electron Affinity– Energy change that occurs when an atom gains an electron. Decreases down a group, increases across a period. Groups Two, Twelve, and Eighteen have full sub-levels and lack affinity. Half-full sub-levels also have low affinity.
  4. Reactivity– In the metals, reactivity increases down a group. Reactivity decreases across a period. In non-metals, it decreases down a group and increases across a period.

Bonding

  • A chemical bond forms when two or more atoms rearrange valence electrons to increase stability.
  • An ionic bond forms when valence electrons are transferred from one atom to another.
  • In ionic compounds, the ions are arranged in a crystal lattice. Strong forces hold the ions together.

Properties of Ionic Compounds:

  • High melting and boiling points.
  • Hard- not easily crushed.
  • Conducts electricity when melted or dissolved because the ions are free to move.
  • In a covalent bond, electrons are shared, forming molecules.
  • Covalent compounds have weaker forces holding the molecules together.

Properties of Covalent Compounds:

  • Lower melting and boiling points.
  • Many covalent compounds are volatile liquids or gases.
  • Softer- easier to crush.
  • Are not conducts of electricity.
  • Since oxygen has a higher electronegativity than hydrogen, oxygen holds onto shared electrons more, giving oxygen a partial negative charge and the hydrogen a partial positive charge.
  • In polar covalent bonds, electrons are shared unequally, creating partially charged ends or poles.
  • In non-polar covalent bonds, electrons are shared equally because atoms have the same electronegativites.
Electronegativity DifferenceType of Bond
Greater than or equal to 1.7Ionic
Between 1.7 and 0.3Polar Covalent
Less than or equal to 0.3Non-Polar Covalent
  • In a metallic bond, electrons are delocalized. (Creates a “sea of electrons”0
  • Metals have luster, are excellent conductors, malleable, and ductile.
  • Intrabonds are within a molecule. (Ionic and covalent)
  • Interbonds are between molecules/ionic units.
  • Ionic takes; covalent shares.

Intermolecular Bonding:

  • Includes dipole-dipole and dipole-hydrogen bonds.
  • Hydrogen bonding is between hydrogen and fluorine, oxygen, or nitrogen.
  • A dipole is a partial positive attached to a partial negative of a different molecule.
  • Melting and boiling points indicate strength of interionic (and covalent) bonds.
  • Phase changes break or form inter bonds.
  • When a substance is aqueous, it is dissolved in water.
  • Dissolving pulls molecules apart. It breaks inter and intra bonds. This is called disassociation.
  • Water is the universal solvent. It tears most things into ions. Water cannot pull apart covalent bonds, only ionic.
  • Electrolytes give off ions when dissolved in water and conducts a charge.
  • Inter covalent bonds are weaker than inter ionic which are weaker than intra ionic which are stronger than intra covalent. (Inter covalent→inter ionic→intra ionic→intra covalent)

Lewis Structures:

  1. Determine the central atom. It will be carbon or the atom with the least amount of atoms.
  2. Bond all remaining atoms to the central atom.
  3. Count valence electrons in the molecule.
  4. Satisfy outside atoms first with remaining electrons.
  5. Satisfy central atom.
  6. If you run out of electrons and everyone is not satisfied, make double or triple bonds.
  • Valance Shell Electron Pair Repulsion (VSEPR)= Unshared pairs pushes bonds close together.
  • Metallic bonds have low electronegativity.
  • Lattice energy is the energy required to break one mol of a substance into individual atoms.
  • Lattice typically describes ionic bonds.
  • The smaller the atoms are in bonding, the shorter the distance between bonds. Shorter bonds are stronger.
  • When charges go up, bond strength goes up.
  • Like dissolves like. Polar dissolves polar, non-polar dissolves non-polar.
  • Ionic bonds are extremely polar.
  • Organic molecules are covalent.

Metallic Bonds:

  • In the Electron Sea Model, all metal atoms in a metallic solid contribute their electrons to a “sea”.
  • Delocalized electrons are not held to one atom.
  • Atoms in metallic bonds are cations.
  • An alloy is a mixture of elements with metallic properties.
  • Atoms bond to become more stable and lower their energy.
  • Atoms do not share electrons equally.
  • The inter attractions of molecules determines the states of matter.
  • Bonds can be single, double, or triple. Triple bonds are the strongest, followed by double.
  • Single bonds are called sigma bonds (σ).
  • Double bonds are sigma and pi bonds.
  • A triple bond contains two pi and one sigma. They also have the shortest bond distance.
  • In an endothermic reaction, more energy is required to break existing bonds than is released.
  • In an exothermic reaction, more energy is released during bond formation than is required to start a reaction.
  • A majority of reactions are exothermic.
  • Resonance occurs when multiple Lewis structures can be written for a molecule or ion. It is shown through the use of dotted lines.
  • Atoms can be satisfied with numbers other than eight. Boron can have six. (Contracted octet) Sulfur, phosphorus, and xenon can have more than eight. (Expanded octet)

Bond Angles

  • Any molecule containing only two atoms has a linear shape.
  • The VSEPR theory can be used to predict the shapes of molecules with more than two atoms.

Intermolecular Forces

  • Intermolecular forces are forces of attraction between molecules.
  • They are weaker than covalent and ionic bonds.
  • There are three types of intermolecular forces:
  1. Dipole-Dipole Forces
    • This is the force of attraction between the positive end of one molecule to the negative end of another molecule.
    • It is the strongest of all the intermolecular forces.
  2. Hydrogen Bonding
    • This occurs in molecules with H-O, H-F, and H-N bonds.
    • The large positive charge on hydrogen is attracted to an unshared pair of electrons on a neighbouring molecule.
  3. London Dispersion Forces
    • This is a weak intermolecular force resulting from constant motion of electrons.
    • It is the only type of intermolecular force between non-polar molecules.

Reaction Types (Chapter Nine)

  1. Synthesis: A+B→C (More reactants, fewer products)
  2. Composition: A→B+C (Few reactants, more products)
  3. Single Replacement: A+BC→AC+B
  4. Double Replacement: AB+CD→AD+CB
  5. Combustion: A+→+O (A will always be a compound containing carbon and hydrogen; it can also include oxygen)

Stoichiometry

1 gram=1 mol=(6.022×10^23 atom/molecule/unit)

Chapter Twelve

Gases:

  • Gas particles are smaller than the space between them. They neither attract nor repel due to distance.
  • Particles move in straight lines until acted on by an outside force.
  • In an elastic collision, no kinetic energy is lost.
  • KE=½mv^2
  • Temperature is the measure of average kinetic energy.
  • Gases have a low density due to the amount of space.
  • Diffusion is the movement of one material through another.
  • Effusion is the diffusion of gas through a tiny opening.
  • (Rate A)/(Rate B)= sqrt(Molar Mass B)/(Molar Mass A)
  • Pressure is the force per unit area. (P=Force/Area)
  • Air pressure is less at higher altitudes than sea level.
  • Manometers measure pressure in a closed container.
  • Dalton’s Law of Partial Pressures states that total pressure is equal to the addition of all partial pressures.
  • Gases are less dense, very compressible, very fluid, less viscus, have less surface tension, and are not very adhesive or cohesive.

Liquids:

  • Fluids are considered anything that can flow.
  • Viscosity is the measure of a liquid’s resistance to flow. Viscosity decreases as temperature
  • Surface tension is the energy required to increase the surface area of a liquid. Surfactants lower surface tensions.
  • Liquids are more dense, in-compressible, less fluid, more viscus, have more surface tension, and are more adhesive and cohesive.

Solids:

  • Unit cells are the smallest arrangement of atoms in a crystal lattice that has the same symmetry as the whole crystal.
  • An allotrope is an element that exists in different forms at the same state.
  • Amorphous solids do not have crystals.

Phase Changes:

  • Deposition is going from a gas to a solid.
  • Evaporation is vaporization at the surface.
  • The triple point is the point on a phase diagram that represents the temperature and pressure at which all three phases can coexist.
  • The critical point indicates the temperature and pressure at which a substance cannot be a liquid.

Gas Laws

Boyle’s Law:

P1 V1=P2 V2

  • Pressure and volume can be in any units as long as they are the same.
  • Sides are inversely proportional.

Charles’ Law:

V1/T1 =V2/T2

  • Volume expands when heated.
  • Temperature and volume are directly proportional.

Gay-Lussac’s Law:

P1/T1 =P2/T2

  • Temperature and pressure are directly proportional.

Combined Gas Law:

(P1 V1)/T1 =(P2 V2)/T2

Ideal Gas Law:

PV=nRT (R[which is a constant]=0.0821Latm/(mol K))

  • Absolute zero is the lowest possible theoretical temperature.
  • Temperature, when used in these equations, must always be in Kelvin.
  • Avogadro’s Principle states that equal volumes of gases at the same temperature and pressure contain equal numbers of pressure.
  • One mol of any gas at STP=22.4 liters of that gas.
  • Molar volume is the volume that one mol of a gas occupies at 0℃ and 1 atm.

Gas Stoichiometry

  • Polar gases do not act ideally.
  • Gases are also not ideal at low temperature or high pressure.
  • If conditions are the same (temperature and pressure are constant), mols can be equal to liters.
  • Liquids are denser than gases.

Chapter Fourteen

  • Heterogeneous mixtures include suspension, thixotropic, and colloids. In these mixtures, individual substances remain distinct.
  • Colloid particles do not settle out.
  • Brownian motion is the random, jerky movements of colloids.
  • The Tyndall Effect is the phenomenon where colloid particles scatter light.
  • Miscible refers to two liquids that are soluble in each other.

Percent by Mass=(Mass of Solute)/(Mass of Solution)x100

Percent by Volume=(Volume of Solute)/(Volume of Solution)x100

Molarity(M)=(Moles of Solute)/(Liters of Solution)

Molality(m)=(Moles of Solute)/(Kilograms of Solvent)

Mole Fraction=(Moles of Solute)/(Moles of Solute+Moles of Solvent)

Dilution Equation: M1 V1=M2 V2

Solvation

  • Solvation is the process of surrounding solute particles with solvent particles to form a solution.
  • Electrolytes form ions. Covalent compounds do not form electrolytes.
  • The heat of solution is the overall energy change that occurs during solution formation.
  • Dissolving is endothermic and exothermic. Stirring makes dissolving faster.
  • To supersaturate a solution, add heat. This allows more solute to be dissolved.
  • Gases dissolve more in cold solutions.
  • Henry’s Law: S1/P1 =S2/P2
  • Colligative properties are physical properties of solutions that are affected by the number of particles instead of their identity.
  • Evaporation starts at the surface. Added substances decrease vapor pressure.
  • At sea level, you cannot heat water above 100. Evaporation is a cooling process. Solutes elevate the boiling point.

Boiling Point Elevation:

∆Tb=Kb m(n)

n is the number of particles in the solute.

Freezing Point Depression:

∆Tf=Kf m(n)

Salt Hydrolysis and Buffers

  • Salts are ions.
  • Salts sometimes change the pH of a solution.
  • Strong bases and weak acids produce basic solutions. Weak bases and strong acids produce acidic solutions. This is due to reactions with water.
  • A strong base and a strong acid will create a neutral solution.
  • Buffers are solutions that resist changes in pH.
  • A buffer is comprised of two compounds. They are a weak acid/conjugate base or a weak base/conjugate acid.
  • According to Le Châtelier’s Principle, with acids the equilibrium shifts left and with a base it shifts right.
  • Buffer capacity is the amount or acid or base a buffer solution can absorb without a significant pH change.
  • There are an equal number of moles of acid/base in a buffer. Buffers have different pH ranges.

Thermochemistry

  • Energy is the ability to do work or produce heat.
  • The SI unit for energy is the joule.
  • Chemical potential energy is the energy stored in a substance because of composition.
  • q represents heat.
  • 1000 calories=1 Calorie (Food)=1 Kcal 4.184 J=1 cal
  • Specific heat is the heat required to raise the temperature of one gram of a substance by 1

q=cmT      m=Mass (g)  c=Specific Heat  T=Temperature (℃) q=Energy lost or gained (+=Gained, -=Lost)

  • Calorimeters determine specific heat.
  • The universe is a system plus its surroundings. The system is the place that contains reactions. Surroundings are everything else in the universe.
  • Enthalpy (H) is the heat content of a system at constant pressure.
  • Enthalpy of Reaction (∆Hrxn) ) is the change in enthalpy.
  • The Thermochemical Equation includes physical states and energy change.
  • Enthalpy of Combustion (∆H_comb) is the enthalpy change for the complete burning of one mol of the substance.
  • Molar Enthalpy of Vaporization (∆H_vap) is the heat required to vaporize one mol of a liquid.
  • q=∆H_vap (n) n=Number of moles
  • Molar Enthalpy of Fusion (∆H_fus) is the heat required to melt one mol of a substance.
  • q=∆H_fus (n)
  • Hess’s Law states that in adding two or more thermochemical equations to produce a final equation, the sum of the enthalpy changes for the individual reactions is the enthalpy change for the final reaction.
  • Standard Enthalpy of Formation (∆H℉) is the change in enthalpy that accompanies the formation of one mol of compound in its standard state.

Reaction Spontaneity

  • Entropy (S) is the measure of the number of possible ways the energy of a system can be distributed. (Disorder)
  • The Second Law of Thermodynamics states that spontaneous processes always proceed to cause the universe’s entropy to increase.
  • Products typically have a higher entropy.

Entropy Rules:

  1. Entropy changes during state changes. (∆Ssystem>0)
  2. Dissolving a gas into a solvent always results in a decrease in entropy. (∆Ssystem<0)
  3. Assuming no physical state change occurs, entropy usually increases when the number of gaseous product particles is greater than the number of gaseous reactant particles. (∆Ssystem>0)
  4. With some exceptions, entropy increases when a solid or liquid dissolves in a solvent. (∆Ssystem>0)
  5. Random motion of particles increase as temperature increases. (∆Ssystem>0)
  • Free energy is energy available to do work.

       ∆Gsystem=∆Hsystem-T∆Ssystem   T=Kelvin    G=J or KJ

  • If is negative, the reaction is spontaneous.
  • If is positive, the reaction is non-spontaneous.

Solubility Rules

  1. All common salts of group one elements and ammonium ions are soluble.
  2. All acetates and nitrates are soluble.
  3. All binary compounds of group 17 elements (except fluoride) with metals are soluble except those of silver, mercury, and lead.
  4. All sulfates are soluble except those of barium, strontium, lead, calcium, silver, and mercury.
  5. Except for those in rule one, carbonates, hydroxides, oxides, sulfides, and phosphates are insoluble.

Redox Reactions

  • In a redox reaction, electrons are exchanged from one atom to another.
  • Oxidation is the loss of electrons; reduction is the gain of electrons.
  • What is reduced is called the “oxidizing agent”; what is oxidized is called the “reducing agent”.

Redox Reaction Balancing Steps:

  1. Divide the equation into two half-reactions, one for oxidation and the other for reduction.
  2. Balance each half-reaction.
    1. First, balance the elements other than H and O.
    2. Next, balance the O atoms by adding as needed.
    3. Then, balance the H atoms by adding as needed.
    4. Finally, balance the charge by adding as needed.
  3. Multiply the half-reactions by integers, if necessary, so that the number of electrons lost in one half-reaction equals the number of electrons gained in the other.
  4. Add the two half-reactions and, if possible, simplify by canceling species appearing on both sides of the combined equation.
  5. Check to make sure that atoms and charges are balanced.

Mass Defect and Binding Energy

Proton Mass=1.00727647 amu                   Neutron Mass=1.006649 amu

  • To find predicted mass, multiply the number of protons and neutrons by their mass. Then, add those numbers together. This will give you the predicted mass.
  • Mass defect is the energy released from nucleus formation. To find it, subtract the actual mass from the predicted mass. (Predicted-Actual=Mass Defect)
  • To find binding energy, convert amu to kilograms. (1 amu=1.66054xkg) Then, plug the numbers into E=mc^2 . (c [speed of light]=2.99792x) Binding energy will be given in joules.

Nuclear Stability

  • Due to the Electro-Static Force, like charges repel.
  • The Nuclear Strong Force holds the nucleus together.
  • When an atom is “unstable”, it is radioactive.
  • Z=Number of Protons  N=Number of Neutrons
  • When Z20, N/Z=1 means that the atom is stable.
  • When Z20, N/Z=1.5 means that the atom is stable.
  • When atomic number is higher than 83, electro-static repulsion causes all atoms to become unstable.

Types of Decay

Alpha                          Beta                            Gamma

  • Decay attempts to stabilize atoms.
  • Alpha particles can be blocked by paper due to their large size.
  • Beta particles can be blocked by metal foil while gamma particles cannot be completely blocked by lead or concrete.

Half-Life and Carbon Dating

  • Half-life is the time it takes for half of an isotope to decay.
  • Carbon dating uses the half-life of a carbon isotope to attempt to discover the age of an artifact.

Nuclear Radiation

  • Penetrating power is the ability of radiation to pass through matter.
  • The band of stability is the area on a graph within which all stable nuclei are found.
  • Elements decay in order to fall within the band of stability.

Positron Emission:

  • This is the decay process that involves the emission of a positron from a nucleus.
  • Positrons have the same mass as an electron but with a positive charge.
  • p→n+e^+ (Proton will become a neutron and positron)

Electron Capture:

  • Occurs when the nucleus draws in a surrounding electron, usually from the lowest energy level.
  • This is the opposite of beta decay.
  • p+e^-→n (Proton and electron becomes a neutron)

Nuclear Reactions

  • Radiochemical dating is the process of determining the age of an object by measuring the amount of an isotope remaining in the object. Decay from U-238 to Pb-206 is used to date rocks.
  • Induced transmutation is artificially creating nuclear reactions by striking nuclei with high-velocity particles.
  • Transuranium elements, elements after uranium, are too unstable to exist in nature.
  • Nuclear fission splits the nucleus into fragments. It can be used to create energy. Fission also leads to chain reactions. (Logrythemic Relationship)
  • When a sample has reached critical mass, the sample is large enough to sustain a chain reaction. Supercritical mass is greater than critical mass.
  • Breeder reactors produce more fuel than they use.
  • Nuclear fusion combines nuclei. Very high temperatures and pressures are needed in order for fusion to take place. Tokamaks create large magnetic fields to contain fusion.
  • Radiotracers are radioisotopes that emit non-ionizing radiation and are used to signal the presence of an element or specific substance.
  • PET scans use positrons to study the brain.

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