Honors Chemistry Semester 1 Final Study Guide

Unit 1 Alchemy

• Scientific Method: a method of procedure that consists of a systematic observation, measurement, experiment, and the formulation and modification of the hypothesis

• Steps of the Scientific Method:

                                                        Observation    ←    ←

                                                             ↓                           ↖

                                                   Hypothesis                  Hypothesis re-stated

                                                              ↓                            (if needed)

                                                        Experiment  → →→↗

                                                      ↙                    ↘

                                        Theory Model                   Law   

                                 ↗                 ↓

Theory Model modified          Prediction

                  (if needed)                   ↓

                                   ↖      Experiment

• Theory: an interpretation or possible explanation of why nature behaves in a particular way; explanation of behavior

• Law: explains why something happened based on the observations, hypotheses, and experiments done; measurable behavior

• Experiment: scientific procedure undertaken to make a discovery, test a hypothesis, or demonstrate a known fact

• Hypothesis: superstition or proposed explanation made on the basis of limited evidence as a starting point for further investigation

• Observation: remark, statement or comment based on something one has seen, heard, or noticed

• Theory Model: a description or representation used to understand the way in which a process works

• Significant Figures Zero’s

  • Sig figs in a measurement included all the digits that can be known precisely plus a last digit that must be estimated

  • Rules

    • Nonzero digits: significant

    • Leading zeros: zeros that precede all the nonzero digits; not significant; aka beginning zeros

    • Captive zeros: zeros that fall between nonzero digits; significant; aka middle zeros

    • Trailing zeros: zeros at the right end of a number; significant only if number is written with a decimal point; aka ending zero

• Measurement and Scientific Notation

  • Number are written as the product of two numbers

    • A coefficient

    • A power of 10 with an exponent

  • Numbers greater than 10 have a positive exponent. The exponent is equal to the number of places that the decimal point has been moved to the left. 

  • Numbers less than 1 have a negative exponent. The exponent is equal to the number of places that the decimal point has been moved to the right. 

  • Number between 10 and 1 don’t really need scientific notation. 

• Defining Matter

  • When adding or subtracting decimals, the answer must have the same number of digits to the right of the decimal point as there are in the measurement having the fewest digits to the right of that decimal point. 

  • When multiplying or dividing decimals, the final answer must contain no more sig figs than the measurement with the least number of sig figs. The position of the decimal is irrelevant. 

• Dimensional Analysis and SI Units

  • The standards are object or natural phenomena that are of constant value, easy to preserve, and reproduce, and practical in size. 

  • Base Units

  • SI units responsible to know

• Volume: SI unit is cubic meter, m³; conversion: 1 cm³=1 mL

• Density: ratio of mass to volume; mass/volume; g/cm³ or g/mL

• Dimensional Analysis

  • Begin with the end in mind (what you’re solving for)

  • List your given

  • Determine the conversion factors from the SI units. You may have more than 1 conversion factor. 

  • Complete your conversion(s)

• Defining Matter

  • Matter: anything that has mass and takes up space. If not matter, it is energy

  • Density: ratio of mass to volume, or mass divided by volume

Unit 2: Basic Building Blocks

• Building Block Terms

  • Matter: anything that has mass and takes up space

  • Substance: particular kind of matter that has a uniform and definite composition ex) sugar, water

  • Element: substance with one type of atom, simplest form of matter, not separated

  • Atoms: fundamental units of elements, not all the same, 100s of different atoms

  • Compound: substance that contains 2 or more elements chemically combined, can be separated, different atoms

  • Chemical formulas: set of symbols a chemist uses to represent a compound

  • Subscripts: (s)=solid, (l)=liquid, (g)=gas, (aq)=aqueous (dissolved in water)

  • Physical Properties: characteristic of a substance that can change without the substance becoming a different substance ex) color, boiling point, melting point, solubility, state of matter, hardness, density, ductility, malleability

  • Physical change: change in atoms or molecules in a substance stays the same, change in appearance not composition ex) metal rusting, glass breaking

  • Chemical Properties: characteristic that describes the ability of a substance to change into a different substance, not reversible ex) rusting, decomposing, flammability, corrosion, reaction to other chemicals

  • Chemical change: forming one or more substances, resulting substance would have a different chemical formula ex) ice melting, paper burning, food digesting

  • Indicators of a chemical change: gas produced, precipitate forms, color change, temperature change, energy change

  • Mixture: blend or two or more substances

  • Homogenous: mixture that is same throughout ex) sugar, water, salt water, solution, wax

  • Heterogenous: mixture containing regions with differing properties ex) blood, eggs, chocolate chip cookies, concrete

  • Law of Conservation of Mass: in any physical or chemical reaction mass is neither created nor destroyed, it is conserved; same mass at beginning and end of reaction

• Periodic Table

• Everything in the Universe:

  • Energy

  • Matter

    • Substances

      • Elements

      • Compounds

      • They can each be seperated by physical

    • Mixtures

      • Homogenous

      • Heterogenous

    • Both Substances & Mixtures can be seperated by Physical Means

• Isotopes and Building Atoms

  • Subatomic Particles

• All about the atom:

  • All neutrons, protons, and electrons are identical except electrons have different energy levels.

  • The nucleus is dense- 99% of the mass of an atom is located in the nucleus

  • The electron cloud is the most dense

  • In a neutral atom, not an ion, the number of electrons is equal to the number of protons

• Atomic Number: the number of protons is always the same as the atomic number, protons define the atom of an element

• Mass Number: equal to the number of protons plus the number of neutrons, electrons have a teeny-tiny mass therefore not included in the mass number

• Isotopes: atom that has the same number of protons but a different number of neutrons

• Atomic mass (weight): weighted average of the mass of the isotopes of an element

  • Find atomic mass by multiplying occurrence percent by atomic mass of isotopes then adding the two products

• Ions: atoms will gain or lose electrons to become stable

  • Cations are positively charged (lose electrons)

  • Anions are negatively charged (gain electrons)

• Dead Chemists

  • John Dalton: atoms of given element are different from those of any other element; atoms of one element can combine with atoms of other elements to form compounds; atoms are indivisible; ancient model

  • JJ Thomson: discovered electrons; atom is divisible; atom is mostly empty space compared to the size of the electron to the size of the atom; Plum Pudding Model

  • Ernest Rutherford: discovered nucleus; atom is mostly empty space; mass is concentrated in a positively charged nucleus (sort of discovered protons); Gold Foil Experiment Model

  • Chadwick: discovered neutrons and isotopes; protons and neutrons in Rutherford’s model; Ray Tubes

  • Niels Bohr: electrons are stationary, they would fall into the positively charged nucleus; planetary model

  • Robert Millikan: discovered charge of electron; oil drop experiment

  • Ernest Schrodinger: calculated the probability of finding an electron in a certain position around the nucleus (energy levels); electron clouds are most dense; quantum mechanical model

Unit 3: Subatomic Particles (Nuclear)

• Sun Formation: the universe was extremely tiny → big bang → atoms formed → galaxy and stars formed due to gravity → cloud of gas and dust formed spinning disk → gas in center collapsed → sun formed

• The Sun and Our Elements

  • Our sun is one of the 100 billion stars in the Milky Way Galaxy; average star (size and mass)

  • Atoms: 91.2% hydrogen and 8.7% helium

  • Mass: 71.0% hydrogen and 27.1% helium

  • As time goes by, the amount of hydrogen will decrease and the amount of helium will increase

  • High temperatures and pressures strip electrons from the atom leaving a positively charged nucleus and free electrons

  • Plasma: mixture of positively charged nuclei with free electrons with little to no order

  • When positively charged nuclei collide, they combine to make a whole new element

  • Four hydrogen nuclei combine to become one helium nuclei=hydrogen fusion

  • Chemical Rxn vs. Nuclear Rxn

• Gravitational equilibrium: outward pressure of nuclear fusion is balanced by the inward pull of gravity; star spends most of its life with these two forces balances

• How energy reaches earth: particles of light called photons carry energy → photons collide over and over again taking 100,000s of years to move to sun’s surface → after reaching the surface, they can move unrestricted at the speed of light to reach earth in 8 mins 20 secs

• Reactions and Radiation

  • Radiation: penetrating rays and particles emitted by a radioactive source

  • Nuclear forces: nuclear reactions involve the nucleus → nucleus opens and protons and neutrons are rearranged (requires a lot of energy); “normal chemical reactions involved electrons

    • Nuclear forces: short range forces that hold the nuclear particles together

  • Chemical reactions: atoms tend to attain stable electron configurations by losing or sharing electrons

  • Nuclear reactions: nuclei of unstable isotopes (radioisotopes) gain stability by undergoing changes by becoming different elements

  • Nuclear binding energy: energy released when a nucleus is formed (or energy required to break apart a nucleus); the higher the binding energy that more tightly they are held together

  • Unstable nuclei want to be stable

    • Undergo changes to their number of protons or neutron to find their stability

  • Types of Radiation

    • Alpha: contain two protons and two neutrons and have a double positive charge

      • Particle Symbol:  ⁴₂He

    • Beta: electron resulting from the breaking apart of a neutron

      • Particle Symbol:  ⁰_-1e

    • Positron: particle that has the same mass as an electron, but has a positive charge and is emitted from the nucleus

      • Particle Symbol:  ⁰₊₁e

    • Electron Capture: an inner orbital electron is captured by the nucleus of its own atom; combines with a proton and neutron is formed

      • Particle Symbol:  ⁰_-1e (reactant side)

    • Gamma Radiation: high energy photon (electromagnetic) emitted by a radioisotope

      • Particle Symbol:  ⁰₀𝛾

  • Properties of Radiation

• Half Life

  • Nuclear stability and decay

    • Nuclear force: attractive force that acts between all nuclear particles that are extremely close together, such as protons and neutrons in a nucleus

      • Dominate over electromagnetic repulsions

    • Band of stability: stable nuclei that do not change over time

  • Half Life: time required for one half of the nuclei of a radioisotope sample to decay to products

    • After each half life, half of the existing radioactive atoms have decayed into atoms of a new element

  • Solve by: list givens (half-life, total time given, initial mass of isotope), determine # of half-lives in the total time given, multiply the mass of the isotope by one-half for each half-life determined in step 2

• Fission and Fusion 

  • Nuclear chain reaction: nuclei of certain isotopes are bombarded with neutrons, they undergo fission, the splitting of a nucleus into smaller fragments

    • Some of the neutrons produced react with other fissionable atoms, producing more neutrons which react with still more fissionable atoms

    • Controlling: nuclear moderation and absorption

  • Nuclear waste: water cools the spent rods and also acts as a radiation shield to reduce radiation levels

  • Nuclear Fusion: nuclei combine to produce nucleus of greater mass

    • Solar: hydrogen nuclei (protons) fuse to make helium nuclei and two positrons

    • Inexpensive and readily available, high temps needed to initiate

  • Fusion reactions , in which small nuclei combine, release much more energy than fission reactions, in which large nuclei split

  • Detecting radiation: geiger counters, scintillation counters, film badges

  • Uses for radioactive material: diagnose medical problems, carbon dating, smoke detectors, x-rays, medical treatment

Unit 4: A Particulate World- Electron Configuration

• Bohr Model

  • Bohr Model

    • Energy levels of an electron is analogous to the rungs of a ladder

    • The electron cannot exist between energy levels, just like you can’t stand between rungs on a ladder

    • A quantum of energy is the amount of energy required to move an electron from one energy level to another 

  • Understanding Electrons Using The Bohr Model

    • Electrons can be found in different shells around the nucleus 

    • Correspond to regions in space that electrons can occupy

      • Like rugs of a ladder, electron can’t be located between the shells 

      • Each shell can only hold a certain number of electrons

    • When full, electrons must go to a  new shell 

    • Electron shells are represented by letter, n (quantum number)

    • Each shell can hold 2n²

  • Understanding Electrons 

    • Understanding Electrons - electrons occupying the outermost shell 

    • Core electrons - electrons located in all of the inner shells

    • Total electrons minus valence electrons equals the number of core electrons in an atom

  • Quantum Mechanical Model - Erwin Schrodinger (1926)

    • Equation for the probability of a single electron being found along a single axis (a axis)

    • The quantum mechanical model is a mathematical solution

    • Has energy levels for electrons

    • Orbits are not circular 

    • It can only tell us the probability of finding an electron a certain distance from the nucleus 

    • Maximum number of electrons that can fit in an energy level is: 2n²

  • 3 things are shown in a electron configuration: 

    • Principal Quantum Number (energy level) or shell (n) - distance from the nucleus

    • Energy subshell or sublevel (Orbital / angular quantum number, 1) - tells you the type shape of the orbital 

      • Electron cloud shape

      • S, p, d, f

    • Number of electrons

    • n= principal quantum number, shell, energy level

      • Energy sublevel (s, p, d, f)

        • Each sublevel has an orbital

          • Each orbital holds 2 electrons each

  • Electron Configuration

• Step 1: figure out electron configuration for element

• Step 2: write noble gas before it [7, then finish the rest of the configuration

• Ions are atoms that have gained or lost electrons to try to be like noble gas. 

  • Cations have a positive charge (+) and have LOST electrons (Li^+)

  • Anions have a negative charge (-) and have GAINED electrons (Cl^-)

• Orbital Diagrams

  • Three Rules for Writing Orbital Diagrams 

    • The Aufbau 

    • The Pauli exclusion principle

    • Hund’s rule

  • The Aufbau Principle: 

    • Electrons enter orbitals of lowest energy first 

      • Orbital are represented by boxes

      • Each orbital holds 2 electrons

      • Within a principal energy level (n), the s is always the lowest energy sublevel 

      • As the principal energy number (n) increases, sublevels begin to overlap. For instance:

        • The 4s is lower in energy than the 3d

        • The 4f is lower in energy than the 5d

  • Hund’s Rule:

    • When electrons  occupy orbitals of equal energy, they don’t pair up until they have to 

  • The Pauli Exclusion Principle: 

    • An atomic orbital may describe at most two electrons

    • To occupy the same orbital electrons must have opposite spins

  • Explanation of Atomic Spectra: 

  • When we write electron configurations , we are writing the lowest energy

  • The energy level, and where the electron starts from, is called its ground state - the lowest energy level

  • Heat, electricity or light can move the electron up to different energy levels. The electrons is now said to be “excited”

  • As the electron falls back to the ground state, it gives the energy back as light

  • The light is color. We see this color at a specific wavelength in the visible spectrum

  • From the wavelength we can use the equation c=λv  we to calculate the frequency at which this occurs

    • c= speed of light which is 2.998 x 10⁸ m/s

    • Using the planck’s constant (n= 6.625 x 10^-34 Js   we can calculate energy need to excite that electron to the next energy level. E=hv where v is the frequency

• Properties of Light Relationship

  • Relationship of two equations:

    • v=c/λ and E= hv

    • Therefore: E= hc/λ

    • h= 6.626 x 10^-34 Js

    • λ= wavelength (nm)-need to convert to m

      • 1 nm=1.0 x 10^-9m

    • c= 2.998 x 10⁸ m/s

Unit 5: Building with Matter

• Intro to Bonding and Naming

  • Ions: atom that has gained or lost an electron

    • Positive ions: cation- smaller than neutral atom

    • Negative ions: anion- larger than neutral atom

  • Bonding: A chemical bond is an attraction between two atoms. The bond is to achieve a more stable state (lower energy state)

  • Chemical Bonds

    • Ionic Bonds (solids)

      • Transfer (exchange electrons); metal and nonmetal; formula units; referred to as salts

      • Formed by an electrostatic attraction between positive (cations) and negative (anions)

      • Naming

        • Start with name of first ion (cation) in the compound

        • Take next ion (anion) in the compound and replace its ending with an “ide” suffix

    • Covalent Bonds (liquids)

      • Share electrons; nonmetal and nonmetal; molecule

      • Molecular covalent- polar covalent ex) ethanol (solids, liquids)

      • Covalent network- nonpolar covalent ex) diamonds (solids)

    • Metallic bonds

      • Sea of electrons; metal and metal; alloy

• Bond Energy, Naming with Polyatomic Ions and Transition Metals

  • Bond Energy:

    • Bonds do not break and form spontaneously- an energy change is required

    • The energy input required to break a bond is known as bond energy

    • Bond energy is important in describing the structure and characteristics of a molecule

    • Used to determine which Lewis Dot structure is most suitable

    • When a bond is strong, there is a higher bond energy because it takes more energy to break a strong bond

    • When the bond order is higher, bond length is shorter; the shorter the bond length means a greater bond energy because of increased electrostatic attraction

  • Ionic Naming and Bonding with Polyatomic Ions

    • Use parentheses to indicate more than one polyatomic ion

    • “Ate” and “ite” indicate polyatomic ions

    • Use criss-cross and reduce method

  • Naming Ionic Compounds Containing Transition Metals

    • Have more than one oxidation state

    • Roman Numerals to indicate the charge of the transition element

    • If a cation is a transition metal (Sn/Pb) then you must always use a roman numeral in parentheses to indicate charge

Chapter 6: Smells (Covalent Bonding)

• Picturing Molecules

  • Empirical Formula: the formula of a compound expressing the smallest whole number ratio of atoms in a compound (all ionic compounds are empirical formula)

  • Molecular Formula: the chemical formula of a molecular substance; tells the number and kind of each atom in a single molecule of a substance; shows the types of atoms in each molecule and the ratios of those atoms to one another

  • Rules for naming covalent compounds

    • Rules for the first element

      • Named just like it is on the periodic table

      • If the molecule has more than one, use a prefix to say how many

    • Rules for naming element

      • End name with an -ide

      • Use a prefix to say how many there are 

• Hydrocarbons and Hybridization

  • Organic Chemistry: the study of compounds contain the element of carbon

  • Carbon:  it has four valence electrons and would like four more electrons to form an octet; single, double, and triple bonds

  • Hybridization: process in which atomic orbitals are mixed to form new additional orbits

    • Sp³ hybridization: 4 single, sigma bonds

    • Sp² hybridization: 2 single bonds and one double bond

    • Sp hybridization: 1 single bond and 1 triple bond

  • Hydrocarbon Functional Groups

    • Only contain hydrogen and carbon

    • Hydrocarbon functional groups include alkanes, alkenes, alkynes, and aromatics

  • Alkanes: saturated hydrocarbon-only single bonds; simplest functional group; general formula: C_nH_n+2; nonpolar

    • 1 carbon- methane

    • 2 carbons- ethane

    • 3 carbons- propane

    • 4 carbons- butane

    • 5 carbons- peptane

    • 6 carbons- hexane

  • Alkenes: hydrocarbons that contain a double covalent bond; general formula: C_nH_n; nonpolar

  • Alkynes: hydrocarbons that contain a triple covalent bond; general formula: C_nH_n-2; nonpolar

  • Aromatic hydrocarbons: hydrocarbons that have six-membered and delocalized electrons; benzene

• Functional Groups

  • Organic molecules have two parts:

    • A carbon backbone which is relatively inert (stable template for functional groups)

    • One or more functional groups

  • A functional group is a set of atoms bonded together in a specific way

  • Functional groups largely define the chemical and physical properties of the compound

  • Functional Groups

    • Alkyl Halide: -halogen

    • Alcohol (camphor): -OH

    • Ether: -O-

    • Amines (fishy): N with room for 3 bonds

    • Ketone (minty): C double bonded to O with room for 2 more bonds

    • Aldehyde (spicy): C double bonded to O and H with room for 1 more bond

    • Carboxylic Acid (putrid): C double bonded to O and bonded to OH and room for one more bond

    • Ester: C double bonded to O single bonded to an O bonded to a C

• HONC 1234

  • H: makes one bond (single)

  • O: makes two bonds (single, double)

  • N: makes three bonds (single, triple)

  • C: makes four bonds (single, double, triple)

• Shapes

  • Linear

  • Bent

  • Tetrahedral

  • Trigonal Pyramidal

  • Octahedral

  • Trigonal Planer

• Polarity

  • Some covalent compounds share their electrons equally between atoms and some do not share equally

  • Partial charges: 

    • Molecules that don’t share their electrons have a partial charge

    • These molecules are called polar molecules or they have a “dipole” moment

      • Molecules having no charge are called non polar

  • When atoms in a molecule share electrons equally, the bond is a nonpolar covalent bond

  • When two different atoms are joined by a covalent bond and the bonding electrons are shared unequally, the bond is a polar covalent bond

  • In a polar molecule, one end of the molecule is slightly negative and one end is slightly positive

  • Bonds and the shape of the molecule also determine polarity

• IMF’s 

  • Attraction between molecules

  • Weaker than intramolecular forces

  • Forces that hold solids and liquids together

    • When a substance melts or boils, intermolecular forces are broken

    • When a substance condenses or freezes, intermolecular forces are formed

  • Types of Intermolecular Forces (Van der Waals)

    • Dipole-dipole

      • Forces that exist between neutral polar molecules

      • Medium strength

    • Hydrogen bonds

      • Strongest

      • Special case of dipole-dipole

      • FON: Fluoride, Oxygen, Nitrogen bonding with Hydrogen only

    • London dispersion forces

      • Electrons are always moving around the nucleus of the atom\

      • Nonpolar molecules

      • Weakest bond 

Unit 7: Tracking Toxins

• Balancing Equations

  • Reactants: starting materials of chemical reaction; go through chemical reaction; listed on left

  • →: “yields”; indicates the chemical reaction

  • Products: new substance formed; listed on right

  • Coefficient: number of molecules

  • Subscripts: number of atoms in molecule

  • Balancing equations: 

    • Get an unbalanced (skeleton) equation

    • Draw boxed around the chemical formulas

    • Make an element inventory

    • Update your inventory using coefficients until balanced