Chemistry CST Review “70 things to know” Solutions

  1. An atom is the basic unit of a chemical element. An element cannot be broken down into a simpler substance.
  2. A physical change can be reversed through physical means. Ex: ice melts. The ice is still water, it’s just liquid. To make it into ice again, put it in the freezer. A chemical change cannot be reversed through physical means. For example, if you bind carbon to oxygen to make carbon monoxide, you now have carbon monoxide which is completely different from carbon and oxygen and you can’t get the original elements back unless you add lots of energy to break the new bond. Some things that indicate that a chemical change has occurred include bubbling, color change, odor produced, heat/light/energy produced, or a solid precipitate forms.
  3. The particles in a solid vibrate in place and are close to each other, while the particles in a liquid slide past each other and have more room to move, and the particles in a gas bump into each other and have more space in between them.
  4. A mixture is two or more substances physically combined. It may be heterogeneous (you can see that the components are different) or homogenous (all components of the mixture are evenly blended). A synonym for a homogenous mixture is “solution.” Solutions can consist of particles in the same (e.g. ethanol or water) or different phases (e.g. water and carbon dioxide). Mixtures, including solutions are not pure substances. The only pure substances are elements and compounds.
  5. Metals are on the left, nonmetals on the right, and metalloids are on the stairstep.
  6. Metals conduct electricity, have a high melting/boiling point, lose electrons to form cations, have low electronegativity, low ionization energy, form metallic bonds (electron sea model)

Nonmetals do not conduct electricity, may have a low melting/boiling point (gases), gain electrons to form anions, have higher electronegativity, will form bonds with other nonmetals in which electrons are shared, have higher ionization energy

  1. A hypothesis is an educated prediction made to explain natural events. A hypothesis can become a theory if it is supported time after time by repeated experimentation.
  2. Experiments are used to test hypotheses. Scientists prefer (usually) to conduct their experiments in a laboratory in which they can minimize anything that might mask the variables of interest.
  3. They also like to use a control (a group that does not receive the experimental treatment) to ensure that the variable they manipulate (independent) really does affect the other variable(s) they are looking at (dependent). For example, if you want to test the effects of fertilizer on tomato plants, your experimental group would be the plants that receive the fertilizer and your control group would be the plants that do not receive the fertilizer. This provides a very important comparison. How do you know the fertilizer affects your plants’ growth? Why not the soil? The sunlight? If you can show that the plant with the fertilizer grows more than the plant without the fertilizer (which also received the same soil, same amount of sunlight, etc.) then you can be fairly confident that the fertilizer caused the increase in growth.
  4. A quantity is an amount of something. It consists of two parts: a number and a unit.
  5. A. 0.0101 m3; B. 0.0000154 kL; C. 0.0000174 km
  6. Density is how compacted a substance is

D= mass/ volume

Units: g/cm3 or g/mL

Density of water: 1 g/mL or 1 g/cm3

17.43 g/ 29.84 mL= 0.5841 g/mL

  1. Proton: nucleus, positive charge, 1 amu
  2. Neutron: nucleus, no charge, 1 amu
  3. Electron: orbitals (electron cloud; outer part), negative charge, ~0 amu
  4. Atomic number is the number of protons in an element. It determines the identity of the element. The mass number is the weighted average of all the masses of the isotopes of an element.
  5. Electrons that are in a higher energy level in the Bohr model are further from the nucleus than those at a lower energy level.
  6. Si: 1s22s22p63s23p2
  7. Cl: [Ne] 3s23p5

Equation: f * λ = c

where:

f = frequency in Hertz (Hz = 1/sec)

λ = wavelength in meters (m)

c = the speed of light and is approximately equal to 3*108m/s

Isolate for frequency, convert your wavelength to meters, and plug it in

E=h*f

h is Planck's constant, which is equal to 6.62 x 10-34 J . s.

  1. phosphorus
  2. s- Alkali metals and alkaline earth metals (1 and 2)

p- nonmetals (13-18)

d- transition metals (3-12)

  1. main group elements are all elements that are in the top part of the periodic table; the lanthanides and actinides are not main group elements
  2. A. alkali metals; B. alkaline earth metals; C. halogens; D. noble gases
  3. Atomic radii decrease going right and up; smallest atom is fluorine
  4. The first ionization energy is the energy that it takes to remove (metals) or add (nonmetals) an electron from an atom. Electronegativity describes the tendency of an atom to attract shared electrons.
  5. Atoms become stable when they achieve a full outer shell of electrons (octet rule or duet rule for hydrogen)
  6. In a covalent bond, two nonmetals share electrons. In an ionic bond, electrons are transferred from a metal to a nonmetal.
  7. Chlorine and oxygen form a covalent bond.
  8. Ionic: metal and nonmetal

Covalent: nonmetal and nonmetal

  1. This is the orbital diagram for fluorine (no picture of calcium available):

  1. write the symbols and add valence electrons one by one, add bonds when two lone electrons from two different atoms are close together
  2. VSEPR stands for valence shell electron pair repulsion; basically, electrons do not want to be next to each other since their negative charges repel, thus the bond angles in a compound will adjust in order to allow the electron pairs to stay as far apart as possible
  3. A. planar; B. bent; C. linear; D. trigonal bipyramidal
  4. Dipole-dipole attractions take place when two or more neutral, polar molecules are oriented such that their positive (+) and negative (-) ends are close to each other.

Because of the attraction between unlike charges, this is a fairly strong type of intermolecular force, and molecules held together by dipole-dipole forces tend to be in the solid or liquid state. Also, for molecules that are about the same size and weight, the strength of the dipole-dipole forces increases as the degree of polarity increases. In other words, the more polar a molecule is, the stronger the dipole-dipole forces it will form with itself and other molecules.

One very important and unique case of the dipole-dipole attraction is known as hydrogen bonding. Hydrogen bonds are not true bonds: they’re just strong attractive forces between the hydrogen on one molecule and a highly electronegative atom on a nearby molecule.

Hydrogen bonds most commonly form between hydrogen atoms and fluorine, oxygen, or nitrogen. This type of intermolecular force is responsible for water’s unique characteristics, such as its high specific heat and boiling point temperature—but more about that later.

London forces are relatively weak forces of attraction that exist between nonpolar molecules and noble gas atoms, like argon (a noble gas) and octane (a hydrocarbon; C8H18). These types of attractive forces are caused by a phenomenon known as instantaneous dipole formation. In this process, electron distribution in the individual molecules suddenly becomes asymmetrical, and the newly formed dipoles are now attracted to one another.

The ease with which the electron cloud of an atom can be distorted to become asymmetrical is called the molecule’s polarizability. Think of this as a probability issue. The greater the number of electrons an electron has, the farther they will be from the nucleus, and the greater the chance for them to shift positions within the molecule. This means that larger nonpolar molecules tend to have stronger London dispersion forces. This is evident when you look at the diatomic elements in group 7, the halogens. All of these diatomic elements are nonpolar, covalently bonded molecules. Now, going down the group, fluorine and chlorine are gases, bromine is a liquid, and iodine is a solid! For nonpolar molecules, the farther you go down the group, the stronger the London dispersion forces.

  1. Water has a high specific heat due to its hydrogen bonding.
  2. Carbon dioxide is nonpolar. Although both C-O double bonds are polar due to the electronegativity difference between carbon and oxygen, the molecule itself is not polar. The polar bonds cancel each other out because the molecule is linear. Water is polar because the O-H bonds are both polar and do not cancel out. Since the molecule is bent, there is a slightly positive side (near hydrogens) and a slightly negative side (near oxygen).
  3. A molecular compound has covalent bonds while an ionic compound has ionic bonds.
  4. A. calcium phosphate; B. lead (II) nitrate; C. copper (II) hydroxide; D. silver (II) oxide; E. potassium perchlorate; F. dinitrogen pentoxide; G. silicon dioxide
  5. A. SO2; B. Si3N4; C. FeCl2; D. CaO; F. CuBr2
  6. When you are trying to find the empirical formula of a compound, your ultimate goal is to determine the subscripts. You get the subscripts by finding out how many moles of each element you have. Take each % and turn it into grams/ 100 grams of the compound. From grams, you can use the molar mass of the element to get moles. Do this for each element. Once you have the amount of moles, they are supposed to be the subscripts. However, subscripts cannot be in decimal form. You will need to multiply all the subscripts by the same number to get each to become a whole number. This may take some trial and error. Remember you want the most simplified formula for the empirical formula. So…if you end up with something like Na8S4O12, divide each subscript by 4.
  7. The molecular formula is not always the simplest formula. It represents the actual formula for a molecule. For example, all sugars have the empirical formula CH2O, while glucose, a specific kind of sugar, has the molecular formula C6H12O6. Again, in this case your ultimate goals is to find the subscripts, which are equivalent to moles. Figure out how many moles of N and O you have by multiplying grams by their molar mass (inverted). Once you have the number of moles of each, you have the empirical formula. Find the molar mass of the empirical formula. Take the molar mass of the actual compound (92 g/mol in this case) and divide it by the mass of the empirical formula to get n. MF=n(EF). Take n and distribute it to each subscript in the empirical formula. Then simplify and you have your molecular formula.
  8. A. 2K(s) + Cl2(g) => 2KCl(s) (synthesis); B. 2C2H6(g) + 7O2(g) => 4CO2(g) +6H2O(g) (combustion);C.HgO + heat => Hg(l) + O2(g) (decomposition); D. AgNO3(aq) + NaCl(aq) => AgCl(s) + NaNO3(aq) (double displacement); E. Ca(s) + NiCl2(aq) => CaCl2(aq) + Ni(s) (single replacement)
  9. 45.1 g of Fe3O4 is produced; ~4 grams of H2 is produced; ~83 grams of the of the excess reactant (Fe) is not used
  10. % yield= [actual/theoretical]x100%; find theoretical yield by using stoichiometry to determine how much product you should be able to produce given the amount of your reactant; the actual yield is what you actually get when you do the experiment; you get ~73% yield
  11. P1V1=P2V2; 3.84 atm
  12. convert Celsius to Kelvin (C+273=K); (V1/T1)=(V2/T2); T2=V2/(V1/T1); 198 K
  13. Total pressure= sum of the partial pressures (Dalton’s law)
  14. First convert grams to moles, then you can use PV=nRT (make sure all your units match those in R!)
  15. do stoichiometry, using 22.4 L/ mol (one mole of an ideal gas occupies 22.4 L at STP); 172 l O2 required; 137 L CO2 produced
  16. Use stoichiometry and your balanced equation to determine the amount of moles of hydrogen (H2) produced from the 15.94 g Na. After you know the amount of moles, you can use PV=nRT to solve for the volume
  17. solutes lower the freezing point and raise the boiling point (as long as they are soluble; remember ionic solutes have a greater affect on the colligative properties because they have a higher i value)
  18. Raising the temperature usually helps solids and liquids dissolve faster. However, gases dissolve better when the liquid solvent is colder. For example, when the ocean is colder it holds more dissolved carbon dioxide and when it warms up, it releases more carbon dioxide because the solubility of carbon dioxide in water decreases with increasing temperature.
  19. heat= thermal energy; heat is energy that does not do work; heat flows from an object with a higher temperature to one with a lower temperature
  20. ionic compounds dissociate into cations and anions in water
  21. increasing temperature, stirring, and surface area usually will cause a solid to dissolve faster. If it is insoluble, nothing will make it dissolve.
  22. “like dissolves like” means that polar solvents dissolve polar solutes and nonpolar solvents dissolve nonpolar solutes
  23. a solution is saturated when the solvent cannot hold any more solute at its current temperature
  24. dissolve one mole of NaCl in one liter of water
  25. use the molar mass to find moles from grams, then M= moles/L
  26. molality(m)= moles solute/ kg solvent; find moles, then use the formula
  27. a Bronsted-Lowry acid is a proton donor, a Bronsted-Lowry base is a proton acceptor
  28. 2H2O(l) => H3O+(aq) + OH-(aq)
  29. a neutralization reaction occurs when equal amounts (moles) of a strong acid and a strong base are mixed. The products are salt and water, which are nearly neutral, ~pH 7
  30. [H3O+]=10-pH ; pH + pOH= 14
  31. pH= -log[H3O+], where the brackets indicate molarity
  32. find the amount of moles of LiOH using moles=MxV, then use stoichiometry and the balanced equation to determine the amount of moles of HSO4, then once you have the moles, use M=moles/V
  33. Q=mcΔT; Q= amount of heat energy in J or kJ, m=mass (make up one; how about 5 grams), c= specific heat (look in textbook), ΔT= T final- T initial
  34. You need to use Q=mL, where L is the latent heat of vaporization for water (look up); you cannot use the equation from #69 because a phase change is occurring here (boiling)