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Pre-AP Chemistry NOTES 9-1 VALENCE ELECTRONS and the FORMATION of IONS

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Pre-AP Chemistry NOTES 9-1 VALENCE ELECTRONS and the FORMATION of IONS Pre-AP Chemistry NOTES 9-1 VALENCE ELECTRONS AND THE FORMATION OF IONS There is a driving force among atoms to gain or lose electrons in order to reach a low-energy, stable state. This stable state is obtained when an atom has an electron configuration of s2p6 in its outermost energy level. The atom is then said to have a full outer shell of electrons. For example: Argon – Krypton – Both of these elements belong to what group on the periodic table?__________________________________ Noble gases (with the exception of helium) have a stable electron configuration (s2p6). This type of configuration, therefore, is called a noble gas configuration. Why is helium stable, even though it doesn’t have an s2p6 configuration? Why is it virtually impossible to get a noble gas to react or bond with anything? VALENCE ELECTRONS The total number of electrons present in the final s-p configuration is very important. These are the electrons which are involved when atoms bond with each other. They are known as valence electrons – the electrons in the outermost energy level (s and p sublevels) involved in bonding. The number of valence electrons (for the representative elements) can be determined by noting the number of the group the element is in: Barium has __________________ valence electrons Carbon has __________________ valence electrons Iodine has ___________________ valence electrons A Lewis Dot Structure (also called an “electron dot structure”) is a method used to show the number of valence electrons in an atom: EXAMPLES: sodium beryllium nitrogen oxygen argon helium FORMATION OF IONS Octet Rule – atoms react by changing the number of their electrons so as to acquire the stable electron structure of a noble gas (s2p6 – 8 valence electrons) In doing this the atoms form ions (positively or negatively charged atoms) EXAMPLES: sodium 1s22s22p63s1 lose 1 electron [Na+] sodium ion aluminum fluorine oxygen FORMATION OF IONS number of number of Lewis dot symbol valence electrons lost or charge ion symbol structure electrons gained Ba S N K C Ne He ELECTRON CONFIGURATION OF IONS (Extra Credit) Metals lose electrons to achieve noble gas configuration and become cations (positively-charged ions). Nonmetals usually gain electrons to become stable, forming anions (negatively-charged ions). S2- Na+ Fe3+ Zn2+ Pre-AP Chemistry NOTES 9-2 CHEMICAL PERIODICITY FOCUS ON FLUORINE (Fluorine is the “Rudy” of the periodic table: Fluorine is little but strong) Atomic Radius - the radius of the atom Atomic radius decreases as you move across the periodic table and increases as you move down Which has the larger atomic radius? chlorine sulfur Which has the larger atomic radius? chlorine bromine Ionization Energy - the amount of energy required to remove an electron from a gaseous atom In general, ionization energy increases as you move across the periodic table and decreases as you move down Which has the larger ionization energy? oxygen nitrogen Which has the larger ionization energy? oxygen sulfur Electronegativity-the attraction an atom has for electrons when it is chemically bonded to another atom In general, electronegativity increases as you move across the periodic table and decreases as you move down Which has the larger electronegativity? magnesium sodium Which has the larger electronegativity? potassium sodium Pre-AP Chemistry NOTES 9-3 CHEMICAL BONDING: METALLIC AND IONIC BONDS Metallic Bond – consists of closely-packed cations surrounded by free-floating valence electrons This type of bond is formed between metals only (either pure or alloys): Why are metals such good conductors of electricity? Why are metals ductile (capable of being drawn into a thin wire) and malleable (capable of being hammered into thin sheets)? Ionic Bond – a bond formed when electrons are transferred between atoms with largely different electronegativity values Sodium wants to lose and electron, while chlorine wants to gain an electron: When the electron is transferred, oppositely-charged ions are formed, which are then attracted to each other by electrostatic forces: The ions are then held together in lattice structures by electromagnetic attractive forces: Ionic compounds do not have molecules CHARACTERISTICS OF IONIC COMPOUNDS: *Most are crystalline solids with high melting points *Many are soluble in polar solvents (such as water) and insoluble in nonpolar solvents *The molten form of the compound conducts electricity (it contains mobile charged particles) – however solids cannot conduct electricity (the charged particles cannot move) *The aqueous solutions of ionic compounds also conduct electricity well (mobile charged particles) *They tend to be brittle and shatter when hammered Pre-AP Chemistry NOTES 9-4 CHEMICAL BONDING: COVALENT BONDS Covalent Bond – a bond formed when two atoms of similar electronegativity values (usually nonmetals) share one or more pairs of electrons. Neither atom will give up an electron, so they share: The two shared electrons are attracted to orbitals in both fluorine atoms and this holds the atoms together. The atoms do this to attain noble gas configuration: Covalent compounds are made up of discrete molecules, rather than lattice structures. Lewis Dot Structures can be used to represent the covalent compound formed: *The “dash” refers to a shared pair of electrons Double Covalent Bond (Double Bond) – made up of 2 shared pairs of electrons Triple Covalent Bond (Triple Bond) – made up of 3 shared pairs of electrons LEWIS STRUCTURES OF MOLECULES: The structure of a molecule can be determined using the Valence Shell Electron Pair Repulsion Theory (VSEPR). The electron geometry of the molecule can be determined using the following guidelines: 1. Add the total number of valence electrons for all atoms in the molecular formula to determine the number of pairs of electrons. (If the species is charged, add one electron for each negative charge and subtract one electron for each positive charge.) 2. Arrange the peripheral atoms around the central atom, placing one electron pair (represented by a “dash”) between each. Subtract the total number of electrons used from the original total. 3. Use the remaining electrons to fulfill the octet rule (eight electrons in the outer shell) on each peripheral atom. (Don’t fill octets for Groups IA, IIA, or IIIA – these don’t get octets.) Subtract the electron pairs used. 4. If any electron pairs still remain, place them on the central atom. 5. Check to be sure that all atoms fulfill their octet. If some atoms do not, then take some of the electrons from the peripheral atoms and form multiple bonds with the central atom. (NOTE: Only C,N,O,P,S form multiple bonds.) 6. Elements with atomic numbers lower than carbon are full with few electrons (ie. beryllium is full with 2 pairs of electrons, boron is full with 3 pairs of electrons, etc.) 7. Nonmetals from VA on up (beginning with phosphorus) can have expanded octets (more than 4 pairs of electrons around the central atom. The “d” orbitals in these atoms are available to form these types of structures. Without the “d” orbitals, atoms are too small to fit more than 4 electron pairs around them. EXAMPLES: CCl4 PH3 - H2O ClO3 + NH4 NaCN - CO2 NO3 Watch for exceptions to the rule: BF3 PCl5 SF6 - ICl2 XeF4 Pre-AP Chemistry NOTES 9-5 POLARITY OF BONDS AND MOLECULES BOND POLARITY A polar bond forms when two atoms of different electronegativity values bond together covalently. This results in pair of electrons that is shared “un-equally” between the two atoms. The diagram below represents a hydrogen atom (on the left) bonded with a fluorine atom (on the right): hydrogen fluorine Since fluorine in the above example has a higher electronegativity value (3.98) than hydrogen (2.20), their shared electron pair is displaced more toward fluorine, resulting in a partial negative “pole” on the fluorine end and a partial positive pole on the hydrogen end – hence, a polar bond. Molecules may be polar or non-polar, depending upon their molecular structure. Water, for instance, has an asymmetrical structure that consists of two hydrogen atoms and two “lone pairs” of electrons in a tetrahedral orientation around a single oxygen atom. This gives the water molecule an uneven distribution of electric charge, resulting in a molecule that has a partial positive pole and a partial negative pole – a polar molecule: Molecules like carbon tetrachloride consist of polar bonds. However, the symmetrical structure of the molecule results in an even distribution of charge – the molecule as a whole is nonpolar in nature: In general, molecules are nonpolar if *there are no lone pairs on the central atom and *all peripheral atoms are the same ELECTRONEGATIVITY AND TYPES OF CHEMICAL BONDING The following scale shows that the type of chemical bonding is determined by electronegativity differences: Electronegativity Difference Type of Bond 0 – 0.39 nonpolar covalent 0.40 – 0.99 moderately polar covalent 1.00 – 1.99 highly polar covalent > 2.00 ionic EXAMPLES: molecule electronegativity difference type of bond HF NaCl Br2 NO Pre-AP Chemistry NOTES 9-6 MOLECULAR GEOMETRY TYPES OF COVALENT BONDS *Single Bond – consists of 1 pair of shared electrons forming a sigma bond (σ) (2 “p” orbitals overlap “head-to-head”) *Double Bond – consists of 2 pairs of shared electrons *1 sigma bond *1 pi bond (π) – formed when 2 “p” orbitals overlap “side-to-side” *Triple Bond – consists of 3 pairs of shared electrons *1 sigma bond 2 pi bonds RELATED FACTS CONCERNING COVALENT BOND TYPES 1. A single sigma bond is stronger than a pi bond 2. The triple bond is the shortest of the bonds and is stronger than the rest 3. The single bond is the longest of the bonds and is weaker than the rest *A single bond has a bond order of “1” *A double bond has a bond order of “2” *A triple bond has a bond order of “3” 4.
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