Chemistry: Chapter 6 - Chemical Bonding
Induced dipoles (6.5)
Electrons of a non polar molecule is temporarily attracted by a polar molecule, because the charged molecule distributes the arrangement of electrons in the non polar molecule, very weak.
Define lattice energy and explain its significance (6.3)
It's the energy released when one mole of an ionic crystalline compound is formed from gaseous ions. Represented in NEGATIVE energy values, it indicates the amount of energy released when crystal lattices form.
List the six basic steps in writing Lewis structures (6.2)
1.) determine the number of atoms in each element. 2.) write the electron dot notation for each type of atom in the molecule. 3.) determine the total number of valence electrons 4.) arrange the atoms to form a skeleton structure, least electronegative atom is central. 5.) count electrons to make sure it equals. 6.) show all lone pairs and resonance structures
Define molecule and molecular formula (6.2)
A molecule is neutral group of atoms held together by covalent bonds. A molecular formula shows the types and numbers of atoms combined in a single molecule of a molecular compound, such as H2O (type: hydrogen and oxygen, number: 2 H, 1 O).
Define a chemical bond (6.1)
A mutual electrical attraction between the NUCLEI and VALENCE ELECTRONS of different of different atoms, which binds the atoms together
Describe the electron-sea model of metallic bonding, and explain why metals are good electrical conductors. (6.4)
A sea of electrons is a lattice of electrons packed together, metals (especially d-block metals) have many vacant outer shells, which allows the orbitals to overlap and electrons to move about, as metals also don't hold on to their valence electrons very strongly. Therefore, due to their highly mobile valence electrons, metals are good conductors of electricity.
Describe hydrogen bonding (6.5)
A special TYPE of dipole-dipole force, in which a HYDROGEN atom that is bonded to a highly electronegative atom, and is attracted to an unshared pair of electrons on a molecule containing O, F, or N (sometimes C). Gives the H atom a large positive charge, and comes very close to electrons, explaining high surface tension, and why ice floats on water.
Explain the relationship between potential energy, distance between approaching atoms and bond length (covalent bonding) (6.2)
As atoms near each-other, the charged particles in their valence shells begin to interact. The positive nuclei and the negatively charged electrons in the outer shells are ATTRACTED to one another. The attractive forces will continue to dominate, and the potential energy will DECREASE until a distance (bond length) is reached where the repulsion between the like charges (2 positive nuclei, etc,) equals the the attraction between the opposite charges. Thus a stable molecule forms, any closer the atoms get, the repulsion becomes GREATER than the attraction.
State the octet rule (6.2)
Chemical compounds tend to form so that each atom, by gaining, losing, or sharing electrons, has an octet of electrons in its highest (occupied) energy level.
Explain why metal surfaces are shiny (6.4)
Electrons have a large freedom of motion in metallic compounds. Metals also contain many orbitals separated by small energy differences. Therefore, metals can absorb a wide range of light frequencies, the electrons immediately move to higher energy levels, and then fall back down to achieve stability, emitting the energy back in the form of light. The "DE-EXCITATION" results in the shiny appearance of metals, electrons constantly changing energy levels.
Discuss the arrangements of ions in crystals. (6.3)
In a crystal lattice structure, which are what ionic compounds consist of, ions consists of cations (positive ions) and anions (negative ions). Since these positive and negative charges attract, these ions arrange themselves so they are "touching" ions they are attracted to, forming a repeating pattern of alternating cations and anions that are locked in place, giving ionic substances a rigid quality.
Why do most atoms form chemical bonds? (6.1)
In order to become stable, most atoms are in state of instability seeking stability, (most) want eight valence electrons.
Describe ionic and covalent bonding (6.1)
Ionic bonding involves atoms with largely different electronegativities bonding by the atom with a large electronegativity stealing an electron to become stable, thus resulting in an attraction (bond) between anions and cations. Covalent boding involves atoms with similar electronegativities bonding through the sharing of an electron, neither can "steal" the electron away so they are locked in a bond
List and compare the distinctive properties of ionic and molecular compounds. (6.3)
Ionic compounds: -Crystalline solids, typically brittle -High melting and boiling points -Conducts electricity in liquid state -Many soluble in water, but not in non-polar liquid Molecular compounds: -Gases, liquids, solids (made of molecules) -Low melting and boiling points -Poor electrical conductors in all phases -Many soluble in non-polar liquids, but not water.
Explain why metals are malleable and ductile but ionic-crystalline compounds are not (6.4)
Ionic crystals are hard and brittle, because the ions are bonded in such a way where a slight shift could cause a chain reaction, causing the bond to break (hence: ionic compounds = brittle). However, metallic compounds are malleable and ductile because the metallic bonding is the SAME in any direction of the grid. Atoms can slide past one another without encountering resistance or breaking bonds.
Compare and contrast a formula for a molecular compound with one for an ionic compound (6.3)
Molecular formulas typically consists of only nonmetals, while ionic formula typically consist of metals and nonmetals. You can also look at the differences between electronegativities of the elements in the formula, the greater the difference, the more ionic character.
Explain what determines molecular polarity. (6.5)
Molecular polarity is the uneven distribution of charge in a molecule, which is determined by the electronegativity differences between the atoms. H2 is non polar, HCl is polar, if the differences is greater than 0.4, it's polar (more ionic character).
Why is most chemical bonding neither purely ionic or purely covalent? (6.1)
Most atoms electronegativity differences fall somewhere on the spectrum between covalent and ionic, and therefore most molecules have some charge that allows them to bond with other molecules.
Classify bond type according to electronegativity differences. (6.1)
Non-polar covalent: 0 - 0.3 Polar covalent: 0.3 - 1.7 Ionic: 1.7 - 3.3 (100% ionic character)
Explain why scientists use resonance structures to represent some molecules (6.2)
Scientists use resonance structures to represent ions or molecules that cannot be represented by a single Lewis structure.
Describe dipole-dipole forces (6.5)
Strongest intermolecular (between molecules) force, Dipole-Dipole describes a force between polar molecules, the higher electronegativity (more electrons) is attracted to the lower electronegativity (less electrons) end of another molecule
Explain VSEPR Theory (6.5)
The properties of molecules depend on boding AND molecular geometry, the three-dimensional arrangement of atoms in space. VSEPR theory accounts for molecular bond ANGLES. Valence Shell Electron Pair Repulsion, the lone pairs surrounding an atom causes these atoms to be far apart as possible.
London Dispersion forces (6.5)
The weakest attractive force that results from electron motion, creates instantaneous dipoles as electrons move around the atom, certain probability electrons will collect occasionally, create a denser, negative side, thus creating a dipole. Acts between all atoms and molecules obviously, only force on noble gases, may increase with atoms with more electrons.
Why do Pentane and Hexane evaporate so fast? (6.5)
Their intermolecular forces are weaker, they aren't polar because they're more symmetrical, so the only forces acting on them are London Dispersion. Less energy to break these bonds make them evaporate faster than water.
Explain how to determine Lewis structures for molecules containing single bonds, multiple bonds, or both (6.2)
To determine Lewis structures, you represent bonds as lines, which represent the sharing of two electrons. When an atom has a double bond, show a double line, and so on.
How does bond energy apply? (6.2)
When the two atoms mentioned form their covalent bond, the release energy in order to become parts of a molecule. The amount of energy released EQUALS the difference between the POTENTIAL energy when the atoms are separate, to when they are bonded. The same amount of energy must be added to separate the atoms. Bond energy is therefore the energy required to break a chemical bond.
