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Chemistry Exam Guide
Chemistry Semester 1 Exam Study Guide
| Question | Answer |
|---|---|
| physical property | a property of matter not involving in its manifestation a chemical change |
| examples of physical properties | color, hardness, boiling point, electrical conductivity, etc. |
| How can you tell the difference between an element & a compound? | an element is a substance made of same type of atoms, whereas a compound is made of different elements in definite proportions |
| atom | the smallest unit into which matter can be divided without the release of electrically charged particles |
| molecule | a group of atoms bonded together |
| compound | a chemical substance composed of many identical molecules containing atoms from more than one chemical element held together by chemical bonds; represented by a formula |
| element | a species of atoms that have a given number of protons in their nuclei, including the pure substance consisting only of that species; represented by a symbol. |
| molecules and compounds | very similar, but molecules can have two or more of the same atoms, or different atoms, but compounds must have two or more different types of atoms |
| isotope | two or more types of atoms of the same element that have different masses; all have the same number of protons & electrons, but different numbers of neutrons |
| atomic number | the amount of protons in each atom of that element |
| mass number | the total number of protons & neutrons that make up the nucleus of an isotope |
| formula for average atomic mass | mass x abundance/100 |
| shortcomings of Bohr's model | 1. focuses only on smaller atoms, for example, hydrogen 2. the chemical characteristics & electrons' arrangement or distribution was not provided. |
| principal quantum number | the main energy level (as 'n'), distance from nucleus going up, the energy increases; symbol is "n = 1, 2, 3, 4, etc." |
| angular momentum quantum number | indicates the shape of the orbital |
| n = 1 | l = 0; s-orbital |
| n = 2 | l = 0; l = 1; s & p orbital |
| n = 3 | l = 0; l = 1; l =2; s, p, d orbital |
| n = 4 | l = 0; l = 1; l = 2; l = 3; s, p, d, f orbital |
| magnetic quantum number | indicates the orientation of the orbital around the nucleus |
| How many orientations are possible in each "s, p, d, f" sublevel? | s: 1 orientation p: 3 orientations d: 5 orientations f: 7 orientations |
| spin quantum number | indicates the spin of electrons |
| the possible values of spin quantum number | +1/2 -1/2 |
| Aufbau principle | electrons fill in from lowest to highest energy levels; meaning that an atom with a lot of electrons will have a lot of different layers with different amounts |
| Pauli principle | states that no 2 electrons can have the same quantum numbers |
| significance of the spin quantum number | describes the state of an electron, including its energy orbital shape & orbital orientation |
| which elements are designated as the alkaline earth metals? | Group 2A: Beryllium, Magnesium, Calcium, Strontium, Barium, Radium |
| Which elements are designated as halogens? | Group 17: Fluorine, chlorine, bromine, iodine, astatine, tennessine |
| Which elements are metalloids | Boron, tellurium, arsenic, silicon, polonium, astatine, antimony, & germanium |
| What are the main group elements | the group elements whose lightest members are represented by helium, lithium, beryllium, boron, carbon, nitrogen, oxygen, & fluorine |
| atomic radius | one half the distance between the nuclei of indentical atoms that are bonded together |
| trends in atomic radius | decreases across a period because effective nuclear charge increases as electron shielding remains constant; increases as you go down a group because the energy level increases. |
| How does ionization energies of main group elements vary across a period & down a group | As you go down a group, it decreases because the number of energy levels increases, so the distance increases. As you go across a period, you increase because the number of proteins increases |
| cation | positively charged ion |
| anion | negatively charged ion |
| valence electrons | the outermost electrons in the last shell |
| how do you know how many valence electrons are in an element? | the number of valence electrons is equal to the atom's main group number |
| electronegativity | attraction between a nucleus & a shared pair of electrons |
| Why is fluorine special in the case of electronegativity? | It has 5 electrons in the 2p shell; since it is so close to the ideal electron configuration, the electrons are held very tightly to the nucleus |
| 3 types of chemical bonds | ionic bonding, covalent bonding, & metallic bonding |
| ionic bonding | the electrostatic attraction between cations & anions |
| covalent bonding | electrostatic attraction between a positively charged nucleus & a shared pair of electrons |
| metallic bonding | electrostatic attraction by the positively charged nuclei & a "sea" of delocalized electrons |
| What is the relationship between electronegativity & the ionic character of a bond? | the relationship is the greater the difference in electronegativity, the more ionic the bond will be |
| polar covalent | unequal distribution of electrons partially positive, partially negative; 0.4-1.7 polarity |
| nonpolar covalent | equal distribution of electrons; 0-0.3 polarity |
| How are bond energies & lengths related? | the shorter the bond length, the greater the bond energy |
| ionic compound | a chemical compound composed of ions held together by electrostatic forces termed ionic compounds |
| 3 properties of ionic compounds | high melting/boiling points, brittle solids at room temperature, most are soluble in H2O |
| What specific property of metals accounts for their unusual electrical conductivity? | The freedom of electrons to move in a network of metal atoms accounts for electrical conductivity |
| What properties of metals contribute to their tendency to form metallic bonding? | metals are excellent heat conductors & having high mobile valence electrons |
| properties of covalent compounds | depends on size & polarity of molecule; can be solids, liquids, or gasses at room temp.; tend to have lower melting/boiling points |
| Lewis structure | represents chemical bonds & valence shell electrons in a molecule |
| intermollecular forces | force between 2 molecules |
| properties of metallic compounds | luster, malleable, ductile, thermal conductivity, high boiling/melting points, & conducts electricity because of freely moving particles |
| alloys | mixture of metal + different metal or metal + nonmetal; changes physical properties usually makes metal stronger |
| polarity | determined by difference in electronegativity -0-0.3: nonpolar -0.4-1.7: polar ->1.7: ionic |
| polyatomic ions | covalent molecules that act like ions |
| small nonpolar molecules | weak IMFs |
| large nonpolar molecules | stronger IMFs |
| polar molecules | strongest IMFs |
| H-bonding | forms a special type of dipole-dipole attraction when a hydrogen atom bonded to a strongly electronegative atom exists in the vicinity of another electronegative atom with a lone pair of electrons; H is bound to F, O, or N in a molecule |
| Dipole * Dipole | attractive forces between the positive end of one polar molecule and the negative end of another polar molecule |
| London Dispersion Forces (LDFs) | a temporary attractive force that results when the electrons in 2 adjacent atoms occupy positions that make the atoms form temporary dipoles. bigger molecules = more e- = more polarity = stronger LDFs |
| ion | an atom or molecule with a net electrical charge |
| ionization energy | amount of energy required to remove a valence electron |
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