Chemistry Notes

Chemistry

• Science dealing with the composition and reactions of substances.
• Example: H2O
• Theory of the atom.

Chemistry Objectives

• Conservation of Matter: Matter changes form but is neither created nor destroyed; it comes from somewhere and goes somewhere.
• Elements: Things are made of elements.
• Atom: The smallest part of an element.
• Sub-atomic particles: An atom is made up of 3 sub-atomic particles.
• Atomic number, atomic mass, and isotopes.
• Elements in the human body: Understanding which elements and their symbols make up the human body and where they come from.
• Valence Shell: Atoms want to have a full outermost shell of electrons, a full “valence shell”. The configuration is often (2:8:8).
• Chemical Reactions and Bonds: Atoms without a full valence shell undergo chemical reactions, forming strong chemical bonds.
• Strong Chemical Bonds: Understanding the 2 strong chemical bonds and how they differ.
• Structure/Function: Shapes of atoms determine the shapes of molecules, which determine the shapes of macromolecules.
• Importance of Carbon and Water: Carbon-carbon bonds and water-water interactions are crucial.
• Polarity and Solubility: Polar molecules and salts dissolve in water, while non-polar molecules dissolve in fat. Understanding hydrophilic vs. hydrophobic molecules and why this occurs.
• Water Characteristics: Including hydrogen bonding and pH (hydrogen, hydrogen, protons!).

Vocabulary

• acid, adhesion, anion, atom, atomic mass, atomic number, ball&stick model, base, buffer, cation, cohesion, compound, covalent bond, electron, electron orbital shells, electronegative oxygen, element, heat capacity, hydrogen bond, hydrophilic, hydrophobic, ion, ionic bond, isotope, matter, molecule, neutron, nonpolar, pH, polar, product, proton, salts, solute, solvent, space- filling model, structural formula, substrate, valence shell, ways to represent molecules.

A Brief History of Chemistry

• ~1,000,000 years ago: Use of fire.
• ~3,000 years ago:
  • Metal ores.
  • Pottery and glazes.
  • Beer and wine production.
  • Extracting chemicals from plants for medicine and perfume.
  • Cheese making.
  • Rendering fat into soap.
• ~300 years ago: Studies of magnetism & electricity.
• ~600 BC: Concept of "Primary Elements" - Earth, Wind, Water, Fire.

Matter

• 3 states of matter
  • Element
    • Examples?
    • How many elements exist?
    • How many of these elements are essential to life?

Periodic Table of Elements

• Listed in order of atomic number.
• Organized based on similar characteristics
• Dmitri Mendeleev In 1865 based on 56 known elements
• http://www.radiochemistry.org/periodictable/index.shtml
• http://www.dayah.com/periodic/

Elements in You

• Where do you get these?
• Know these symbols
• The big 6; Carbon, Hydrogen, Nitrogen, Oxygen, Phosphorus, and Sulfur

Atom

• Composed of 3 subatomic particles:
• Bohr Atomic Model- 1913
  • Nucleus (Protons and Neutrons)
  • Orbital shells (electrons)
  • Helium Atom
    • Electrons
    • Protons
    • Neutrons

Characteristics of Subatomic Particles

• Electrons:
  • Symbol: e-
  • Location: Orbitals
  • Charge: -
  • Mass: ~0
  • Function: Bonding
• Neutrons:
  • Symbol: n
  • Location: Nucleus
  • Charge: 0
  • Mass: 1
  • Function: Nuclear stability
• Protons:
  • Symbol: p
  • Location: Nucleus
  • Charge: +
  • Mass: 1
  • Function: Identity
• The number of protons gives an element its unique characteristics.

Atomic Number and Mass

• The number of protons is usually = number of neutrons = number of electrons
• Ex. Carbon (atomic number is 6)
• Atomic number = the number of protons in the nucleus of an atom
• $^{12}C$
• 6 protons
• 6 neutrons
• 6 electrons
• atoms very much do not move like this

Periodic Table Information

• Element name (e.g., Carbon)
• Element symbol (e.g., C)
• Atomic number: The number of protons in the element.
• Atomic mass: The average mass of all isotopes of the element.

Atomic Mass

• Atomic mass: The sum of the mass of all its subatomic particles (the relative weight of isotopes averaged).
• Ex. Carbon 6 electrons + 6 neutrons + 6 protons 0 + 6 + 6 =? Why is carbon’s atomic mass not equal to 12?
• Units of mass (AMU) (Daltons) !
• Atomic mass of all forms of that element vs. of a single atom

Isotopes

• Atoms having the same number of protons but differing numbers of neutrons.
• Ex. Carbon has 3 isotopes ($^{12}C$, $^{13}C$, $^{14}C$)
  • 98.9% is “regular” carbon 12 ($^{12}C$) -> 6 neutrons
  • 1.1% is carbon 13 ($^{13}C$) -> 7 neutrons
  • trace is carbon 14 ($^{14}C$) -> 8 neutrons
• $^{14}C$ is unstable “radioactive”
• There are actually 15 carbon isotopes $\text{ }^{9}C$ – $\text{ }^{22}C$
• All are carbon! Why?

Isotopes and Climate Science

• How does science know the increase in atmospheric CO2 is due to the burning of fossil fuels?
• Isotope ratios
  • Volcanoes -> low $\frac{^{13}C}{^{12}C}$ ratios
  • Volcanoes ~200 million tons CO2/yr.
  • Fossil fuels -> high $\frac{^{13}C}{^{12}C}$ ratios
  • Fossil fuel ~25 billion tons CO2/yr.
• Which is bigger? 200 million or 25 billion?
• April 2025 summit Sundhnúkagígar, Iceland
• CO2 Also Ash & $\text{SO}_2$

Electrons and Reactivity

• Electrons determine the reactivity of an element (chemical bonding).
• “orbital shells” are in specific areas
• What determines if a reaction happens?
• See 2, 8, 8 rule
• Notice e- often hang out as pairs
• Electrons e-

Electron Orbital Shells

• What do these e- orbital shells look like?
• First orbital shell: One spherical orbital (2s)
• Three dumbbell-shaped orbitals (2p)

Electron Shell Capacity

• Max # e-
  • 1st shell: 2
  • 2nd shell: 8
  • 3rd shell: 8
• What happens if a shell is not full?
• What is this element?
• Note: e- in outer orbital shells have more energy than do inner shells.
• Orbital shells “2,8,8” rule or octet rule
• 18p+ e-
• 2 8 8

Elements with Full Outer Shells

• Noble gases have full outer shells.
• These elements have full outer shells
• "noble gases"
• 2
• He

Using the Periodic Table

• Periodic table (The important rows for our class)
• Know the important chemical symbols
• Recognize and be able to use the atomic number to tell:
  • Number of protons
  • Number of electrons and where they would be (e- shells)
• Recognize and be able to use the atomic mass to tell:
  • Number of neutrons, and to figure Molecular mass of molecules.
• 2 8 8

Electron Configuration Examples

• Hydrogen has one electron. Where is the one e-? How many more e- does it want to have?
• Helium has two electrons. Where are the e-? How many more e- does it want to have?
• Max # Hydrogen (H)
  • 1st shell 2 1 Not full
  • 2nd shell 8 -
  • 3rd shell 8 -

Sodium Electron Configuration

• Sodium has 11 protons thus it has 11 electrons Where are these electrons?
• Na atomic # 11
• Max # e- Sodium (Na)
  • 1st Shell 2 2
  • 2nd Shell 8 8
  • 3rd Shell 8 1
• How many more electrons does sodium “want”?

Valence Shell and Reactivity

• If the outermost shell of electrons is not full, the atom is “reactive”. Valence shell
• How many valence shell electrons does sodium have?
• How many valence shell electrons does it want?
• How does it get a full valence shell of electrons?

Strong Chemical Bonds

• Strong chemical bonds in living organisms
  • Ionic Bonds
  • Covalent Bonds
• It is all about the valence shell electrons

Ionic Bonds

• Ionic bonds – form from opposite charge attractions. How form?
• “salts”
• These give up electrons
• These take electrons

Salts

• “Salt” What is salt?
• How is salt different from sugar?
• What do you need to know about salts?
• $CuSO_4$
• Salts- molecules held together by ionic bonds

Example: NaCl (Table Salt)

• Number of protons:
  • In sodium (Na)?
  • In chlorine (Cl)?
• Number of electrons?
• Ionization of sodium and of chlorine atoms

Formation of Ions

• Na loses 1e- (10 e- and 11 protons+) (has one + charge)
• Cl gains 1e- (18 e- and 17 protons +) (has one - charge)
• ${Na}^+$ ${Cl}^-$
• How many e- does Na want?
• How many e- does Cl want?
• 17p+ 17 e- 17p+ 18 e-
• Before e- moved After e- “moved”
• Electrons are fickle
• 11p+ 11e- 11p+ 10e-
• 11p+ 11p+
• 17p+ 17p+

Ions

• Ion- an atom (or molecule) that has lost or gained one or more e-, giving the atom an electrical charge.
• Ex. Sodium ion (1 + charge) ${Na}^+$ (cation)
• Chloride ion (1 - charge) {Cl}^-$ (anion)
• Result- opposite charges attract
• How many e- each have?
• Ion vs. Salt
• 11p+ 10e- 17p+ 18e-
• + -

Ionic Bonds and Salts

• Ionic bond- occurs between two oppositely charged ions
• Salt- Any compound formed from ionic bonds.
• Ion vs. Salt
• NaCl
• +
• -

Salt Crystal Structure

• {Cl}^-
• {Na}^+
• Salts form strong ionic bonds between oppositely charged ions forming crystals
• How do salts dissolve in water?
• How does water disrupt strong ionic bonds?

Key Points About Salts

• “Salt” is not just NaCl
• Salts are molecules held together by ionic bonds
• Ionic bonds are formed when oppositely charged atoms (or molecules) attract one another.
• Ions result from transfer of negatively charged electrons.
• Ions are important in biological systems

Common Ions and Salts

• Some common ions
  • {H}^+
  • {K}^+
  • {Cl}^-
  • {Na}^+
  • {Mg}^{++}
  • {OH}^-
  • NH_4 ^+
  • CO_3^{-2}
• Common salts (NaCl, MgCl2$, KCl,$MgSO4$,$CaCO3$,$NaHCO3 etc.)
• How can you recognize something is an ion?
• {MgSO_4}
• NaHCO_3
• KCl
• “Electrolytes”
• NaCl, KH2PO4$,$Na3C6H5O7

Covalent Bonds

• Strong bonds cont.
• 2. Covalent bonds - form when atoms share electron pairs.
  • strongest type of bond
• Hydrogen also forms covalent bonds
• In this class (H,C,N,O,P,S)

Covalent Bond Example: Hydrogen and Carbon

• Substrates (reactants) -> Product
• One atom of carbon can form 4 covalent bonds (one each to 4 hydrogen atoms)
• This is a chemical reaction
• Carbon + Hydrogen -> Methane

Methane Structure

• What these do these lines mean?
• How many valence e- does H have?
• How many valence e- does C have?
• Name of this molecule?
• Where are carbon’s 8 valence electrons?
• Note 3D shape:
• Tetrahedral
• Common name(s)
• Chemical name
• Non-polar
• Note that the covalent bonds are equally surrounding the carbon. This means methane is non-polar!
• e- e- e-e- e-e- e- e-

Molecular Mass

• Molecular Mass
• Atomic mass can be used to determine molecular mass.
• What is the molecular mass of methane (CH_4)?
• Sum of its components
• Note atomic mass is not shown in the other periodic table I am using in my notes

Conservation of Mass

• Can we turn lead into gold?
• No…. Well…. No easily
• "Immutability of elements'
• Law of the conservation of mass (matter)
• Matter can be rearranged but “matter cannot be created or destroyed”
• Ex. When burning gasoline, octane (before and after)
• octane + oxygen -> carbon dioxide + water
• Note: no elements have changed!
• Where do the chemical elements go?
• Where does the energy in the octane go?
• Where did the energy in octane come from?
• How is this reaction reversed?
• Where lead in PDX school drinking water coming from?
• Arsenic found in drinking water at Victor Point Elementary School in Silverton

Organic Molecules

• “Organic molecules” - contain carbon (and at least one hydrogen)
• Is pure stream water, “organic”?
• What is organic food?
• Is organic food “better”?
• Note: Emergent Properties
• What do these lines represent?

Compounds

• Compound
• A pure substance formed when atoms of different elements form bonds.
• The number of atoms of each element is written as a subscript.
• Examples:
  • CO_2 carbon dioxide
  • H_2O water
  • CH_4 methane
  • C6H{12}O_6 glucose
  • NaCl table salt

Molecules

• Molecule
• One or more atoms joined by covalent bonds. (can be of the same or different element)
• The number of molecules is written as a coefficient.
• Examples:
  • CO_2 1 molecule of carbon dioxide
  • 4CO_2 4 molecules of carbon dioxide
  • 2C6H{12}O_6 2 molecules of glucose
  • 6O_2 6 molecules of oxygen
• Note: oxygen gas is not a compound

Compounds vs. Molecules

• All compounds are molecules, but not all molecules are compounds
• -Is H_2O (water) a compound or a molecule?
• -Is O_2 (oxygen gas) a compound or a molecule?
• Which of these is a chemical?

Covalent Bonds in Water

• Example continued
• covalent bonds in water H_2O
• Bulk biological elements
• Trace elements
• Possibly essential
• trace elements

Water Molecule

• Example covalent bonds in a water molecule H_2O
• How many e- does O want?
• How many e- does H want?

Water Molecule Formation

• O + 2H ® H_2O
• Why are the H atoms at a weird angle in the water molecule?
• Result: Unequal sharing of electrons
• Oxygen very much wants e- (very electonegative)

Polarity of Water

• Water is polar Result
• Part of the water molecule (H side) has a slight positive charge (+)
• Part of the water molecule (O side) has a slight negative charge (-)
• Recall e- often hang out as pairs
• Note shape: tetrahedral
• What does that mean?
• e- e- e- e-e- e- e- e-

Polar vs. Non-Polar Molecules

• Polar molecules- molecules with one part that is slightly positive and another part that is slightly negative
• ex. Water, ammonia, sugars, some amino acids, DNA,
• Non-polar molecules- molecules with no charges
• ex. Methane, fats, gasoline, etc.

Water's Shape & Hydrogen Bonding

• (slightly -)
• H
• H
• (slightly +) (slightly +)
• result of being polar:
  • "The shape of water"
• Space-filling
• model of water
• H-bonding (see below)

Hydrogen Bonding

• Chemical bonds continued (weak chemical bonding)
• 3. Hydrogen bonding (H- bond)- formed when the positive end of one polar molecule, such as water, is attracted to the negative end of another polar molecule.
• Result from slight opposite charge attraction
• These are not really “bonds”
• Do you understand this figure?
• H-bonds

Examples of Hydrogen Bonding

• 3. Hydrogen bonding (weak chemical bonding)
• Ex. H- bonds in H_2O
• Ex. H- bonds in DNA
• Ex. H- bonds in a sugar crystal and as the sugar dissolves into water

Hydrophilic vs. Hydrophobic

• Things in our bodies/cells are either water soluble or fat soluble
• Hydrophilic- Water soluble
• Hydrophobic- Fat soluble (water insoluble)
• Solubility
• Solubility in our bodies and cells
• Why don’t oil and water mix?
• Based on this experiment, what do you hypothesize about the solubility of the red coloring?
• Is the food color hydrophobic or hydrophilic?
• Do aspirin dissolve like this?
• So far:
  • Salts- That dissolve to form ions (ex. {Na}^+$${H}^+$,${K}^+$,${Cl}^-$,${Mg}^{++})
  • Polar things- (ex. water, DNA, sugars, some amino acids, etc.)
  • Non-polar things- (fats, waxes, methane, oils, etc.)

Solvents and Solutes

• Things that dissolve in water Hydrophilic
  • Polar molecules
  • Salts and their charged ions
  • Sugars glucose
• Solvent- a chemical in which other molecules dissolve.
• Solute- a chemical that dissolves.
• ex. When making Kool-Aid solution, Kool-Aid is the solute, water is the solvent
• Note: some gasses dissolve in water (O2$,$CO2)
• Solubility Solutions
• Ions

Hydrophobic Substances

• Things that do not dissolve in water but do dissolve in oils. Hydrophobic
  • Non-Polar/Non-Charged molecules
• Solvent- a chemical in which other molecules dissolve.
• Solute- a chemical that dissolves.
• ex. Chili infused oil, oils from chili are the solutes, vegetable oil is the solvent.
• Note: (hydrocarbons) are non-polar
• Octane (gasoline)
• CH_4 (Methane)
• Solubility Solutions

Cellular Solubility

• Within cells there are hydrophilic regions separate from hydrophobic regions
• Solubility of things in the cell
  • Water soluble things “hydrophilic”
    • Polar things (ex.?)
• Charged things (ex.?)
• Lipid soluble things “hydrophobic”
  • Non-charged, non-polar things (ex.?)
• Where in the cell would you find these things?
• Solubility

Covalent Bond Summary

• Summary. What to know about covalent bonds?
• the # of e- need to fill valence shell = # of covalent bonds
  • Hydrogen, forms 1 covalent bond
  • Carbon, forms 4 covalent bonds
  • Nitrogen, forms 3 covalent bonds
  • Oxygen, forms 2 covalent bonds

Common Molecules

• Some common molecules
• H_2O (know water)
• (you will not be tested on these yet:)
• O_2 (oxygen gas)
• CO_2 (carbon dioxide)
• C6H{12}O_6 (glucose)
• HPO_4 (Phosphate)
• C{738}H{1166}N{812}O{203}S_2Fe (hemoglobin)
• ATP (adenosine tri-phosphate)
• Note ATP does not represent a chemical formula.

Water

• Water Life giving water
• Water exists on Earth in all three forms of matter
• Water is polar
• Water forms H-bonds

Water Characteristics

• Characteristics of water resulting from H-bonding and its polar nature.
  • Cohesion
  • Surface tension
  • Adhesion
  • Is a solvent
  • High heat capacity
  • High heat of vaporization
  • Ice floats
  • Moderate temperature changes

Cohesion

• Characteristics of water resulting from H-bonding and its polar nature.
• 1. Cohesion- water sticks to itself. Why important?
• Ex. Holds water together as it is Sucked up straws.
• Ex. Holds water together as it is Sucked up trees.
• Ex. Allow for surface tension

Surface Tension

• 2. Surface tension- cohesion of water molecules creates a barrier at air interface. Importance.
• Ex. Prevents penetration by this insect

Adhesion and Capillary Action

• 3. Adhesion- water sticks to other things. Coats surfaces (ex. roots and cells) Importance (see below)
• Water is a solvent for
  • polar (hydrophilic) molecules
  • ions
• Adhesion & Cohesion allows:
  • Drop formation
  • Capillary action.
• Ex. Holds water in soils.
• Ex. How sponges hold water NOT blood capillaries

Water as a Solvent for Polar Molecules

• 4. Water is a good solvent for hydrophilic substances
  • polar (hydrophilic) molecules
• How do polar molecules like sugars dissolve. what happens to sugar’s:
  • H- bonds?
  • Covalent bonds?
• Water is the solvent.
• Glucose is the solute.
• Why do cotton balls absorb water?

Water as a Solvent for Ions

• 5. Water is a solvent for hydrophilic substances
  • Ions
• How water dissolves ions
• Ionic bonds broken and replaced by H-bonds
• Note: dissolved ions are essential for many biological functions
• Water is the solvent.
• NaCl is the solute.
• Note: NaCl dissolved in water is no longer “salt”. Why?

Sugar vs. Salt Crystals

• How are sugar crystals different from salt crystals?
• Why is sugar syrupy/sticky?
• Why isn’t salt sticky?
• Dotted lines are hydrogen bonds
• Solid lines are ionic bonds

Why Water Doesn't Dissolve Hydrophobic Things

• Why does water NOT dissolve hydrophobic things?
• Structure/Function Hydrophilic Hydrophobic
• Note : Nothing for water to grab onto
• Glucose \text{C}6H{12}O_6
• Octane \text{C}8H{18}
• Which above is molecules name, molecular formula, structural formula, ball&stick, or space-filling model?

High Heat Capacity of Water

• High heat capacity - H-bonds require energy to break, so water stabilizes temperature. Importance:
• Ex. Minimizes temperature extremes on earth (especially near large bodies of water)
• Ex. Minimizes temperature fluctuations in you.

High Heat of Vaporization

• 6. High heat of vaporization- It takes lots of energy to break all H-bonds forming water vapor Importance:
• Ex. “Evaporative cooling”
• Homeostasis!
• Sweaty Serena Williams
• What happens to covalent bonds?

Ice Floats

• 7. Ice floats- (water is densest at 4°C (39.2°F))
• What would happen if ice did not float?
• Note: 6 water molecules form a hexagon

Water Dissociation and pH

• Water disassociates (form ions) H_2O$${OH}^-$+${H}^+
• pH- The measure of {H}^+ concentration in a solution
• pH= -log [H^+]
• Pure water disassociates to an equal amount of {OH}^- & {H}^+
• concentration of \text{10}^{-7} moles per liter = pH 7
• pH The “Power of H”
• Free {H}^+ ions are called “protons”
• Don’t worry about this equation pH 7 is referred as “neutral pH”
• What is this?

Acids and Bases

• Acids & Bases pH scale is measure of acidity/alkalinity based on {H}^+ concentration.
• Note: pH scale is -logarithmic.
• Ex. pH 5 is ten times more acidic than pH 6
• Acids - add {H}^+ to a solution (i.e. lower pH below 7)
• Ex. Hydrochloric acid
• HCl {H}^+$+${Cl}^-
• Free {H}^+ ions are very reactive
• Figure is a bit confusing blood and tears do not have a pH of 7
• Blood has a pH ~7.4

More on Acids and Bases

• Acids & Bases pH scale is measure of acidity/alkalinity based on {H}^+ concentration.
• Note: pH scale is -logarithmic.
• Ex. pH 6 is 1/10th as acidic than pH 5
• Bases - remove {H}^+ from solution (i.e. raise pH above 7)
• Ex. Lye
• NaOH {Na}^+$+${OH}^-
• Free {OH}^- is very reactive.
• Removes free {H}^+ from solution

Buffers

• Buffers- resist changes in pH
• Ex. Bicarbonate/Carbonic-acid in blood stabilizes pH 7.3-7.5
• Homeostasis!
• If blood become too acidic, bicarbonate accepts H to form carbonic acid
• H2CO3$+${OH}^-$→${HCO_3}^- + (H-O-H)
• {H_2O}
• {H2O}$+$CO2$->$H2CO3 (carbonic acid)
• What does the buffer above buffer at?
• Would this work in the blood?
• {HCO3}^-$+${H}^+$→$H2CO_3

pH and Biological Processes

• At what pH do biological processes occur?
• What is the effect of high pH (alkaline)?
• What is the effect of low pH (acidic)?
• pH has big effect on the shape of proteins! (see macromolecule lecture)
• Neutral pH is does not mean it has neutral effect.
• pH 7 would harm our stomach
• You know these buffers?

Common Buffers

• Some chemicals that act as buffers in the body, helping to maintain pH balance (Homeostasis):
  • Bicarbonate Buffer System ({HCO3}^-$/$H2CO_3)
    • Components: Bicarbonate ion ({HCO3}^-$/ Carbonic acid ($H2CO_3)
    • pH Range: Effective around pH 7.35–7.45 (blood plasma)
    • Function: Primary buffer in blood, regulating pH by neutralizing excess acids or bases.
    • Small Intestine Buffer: Neutralizes stomach acid as chyme enters the small intestine.
  • Phosphate Buffer System ({HPO4}^{-2}$/${H2PO_4}^-
    • Components: Monohydrogen phosphate ({HPO4}^{-2}$/ Dihydrogen phosphate (${H2PO_4}^-$$
    • pH Range: Most effective around pH 6.8 (intracellular fluids and kidneys)
    • Function: Buffers pH changes in urine and intracellular fluid, especially in the kidneys.

Other Atomic Attractive Forces

• van der Waals attractions- attractions that occur within or between molecules. (surface attractions due to minute electro-static charges)
• “Allows gecko lizards to walk up walls”
• Help hold non-polar molecules together (ex. Oil)
• Help maintain shape of LARGE molecules (ex. Proteins)

Review of Chemical Bonds

• Review Chemical bonds
  • I. Strong chemical bonds
    • 1. Ionic bonds- are opposite charge attractions between ions. (electrical)
    • 2. Covalent bonds involve sharing of electrons between atoms.
  • II. Weak chemical bonds
    • 1. H-Bonds link separate polar molecules

Solubility Review

• Amphiphilic- molecules where one part is lipid soluble, and one part is water soluble.
  • Ex. Detergents, Ethanol
  • What about phospholipids?
• Molecules whose solubility depends on pH (pH effects molecules charge)
  • Ex. aspirin has a pKa of 3.5
• Review Solubility: hydrophobic, hydrophilic and………

Chemistry Test Study Guide

• Chemistry test study guide
  • Know the vocabulary
  • Know the difference between a compound and a molecule
  • What is the role of protons? electrons? neutrons?
  • Be able to draw and label an atom. Ex. Si
  • Know the definitions for the chemical bonds
  • Know how to interpret a chemical formula and a chemical equation Ex. O + 2H -> H2O
  • What is a H-bond? What is not a H-bond?
  • Water is polar. Why?
  • Know why some things are hydrophobic vs. hydrophilic
  • How does hydrophobic/