3.1 Test Review
Objective 1: Analyze the relationship between the valence (outermost) electrons of an atom and the type of bond formed between atoms.
1. Understand the relationship between groups or families on the periodic table and the number of valence electrons in atoms of the elements in the group or family. Valence electrons are electrons in the outermost principal energy level (n) of an atom. Be able to qualify groups of elements based on their numbers of valence shell electrons. Remember main groups are only s and p block groups (d and f block groups do not constitute main groups). There are only 8 main groups.
2. Recognize the difference between ionic, covalent, and metallic bonds. Know which bond type is formed when electrons are transferred between a metal and a nonmetal. Know which bond type is formed when electrons are shared between two nonmetals. Know which bond type is formed when valence shell electrons are shared by all metal atoms.
3. Know properties of substances formed by ionic bonds, covalent bonds, and metallic bonds based on the position of electrons in the valence shells of atoms of elements. Properties include melting point, boiling point, lattice energy hardness, brittleness, bond energy, bond length, bond strength, electrical conductivity, thermal conductivity, malleability, ductility, luster, and enthalpy of vaporization.
4. Understand the octet rule. Most chemical compounds form bonds so each atom has eight electrons in each valence shell. Noble gases have eight electrons in the valence shell and full s and p sublevels.
5. Review groupings of atoms on the periodic table like metals, nonmetals, and semimetals or metalloids. Be able to identify symbols and names of common elements within each of the above groupings.
6. Be able to draw electron-dot symbols, orbital notations, electron configurations for elements on the periodic table. Be able to assign quantum numbers (n, l, m, s) to orbitals within atoms on the periodic table.
7. Compare and contrast chemical bonding to nuclear transmutation. Chemical reactions results from the exchange or sharing of electrons in outer shells (valence shells) of atoms. Nuclear reactions results from changes in the nucleus (protons and neutrons).
8. Review electronegativity. What does electronegativity mean and what does it have to do with atoms and electrons? Be able to identify the most electronegative element given a group of elements from the periodic table. Review electronegativity trends. Review other periodic trends including atomic radii, ionic radii, ionization energy, and electron affinity.
9. Remember the difference between ground state and excited state in the Bohr model of the atom. Ground state is the lowest energy state of electrons in atoms. Excited state is electrons in higher energy orbitals than ground state.
10. Be able to predict how atoms in main groups on the periodic table will react upon becoming proximate to each other (close together). Atoms in Main Group I will lose an electron to atoms in Main Group VII. For example, sodium will lose an electron to chlorine. What other combinations are possible between atoms on the periodic table. Could Main Group I react with Main Group VI? In what ratio?
11. Recognize the association between mobile or fluid electrons and electrical conductivity. Why are metals conductive?
12. What is the difference between a cation and an anion? Do metals tend to lose electrons and become cations (positive) or gain electrons and become anions (negative)?
13. Review the definition of an ion. How do ions differ from atoms? What does the symbol of an ion look like? Does an atom lose a significant amount of mass when an ion forms (remember electrons have relatively no mass compared to protons and neutrons). How are ions different than isotopes?
14. Study the properties of hydrogen. Recognize that hydrogen can be a metal (lose electrons) or a nonmetal (gains electrons). Most of the time, hydrogen will be a nonmetal. For this class, consider hydrogen a nonmetal.