Grade 11 Chemistry

Página 1

CHEMICAL BONDING (Unit 4)

ABE naming method:

-A is the number of central atoms

-B is the number of atoms attached to the central atom

-E is the number of lone pairs of electrons on a central atom

-The number of each is written as a subscript

Bond Dissociation Enthalpy: The energy required to break the bonds between one mole of bonded energy. The enthalpy depends on how the molecules are bonded, including both intramolecular and intermolecular forces that can bond molecules.

The Octet Rule:

-a rule of thumb for how elements tend to combine

-generally an atom will be most stable when there are 8 electrons in its valence shell (like a noble gas) and thus it will strive to complete this valence shell when combining with other elements by gaining or losing electrons.

-not all elements are most stable in this form

Covalent Bonding:

-intramolecular force

-non-metals bonding with non-metals

-polar covalent bonding: uneven sharing of electrons that is based on the electronegativity of elements, stronger intermolecular forces than normal covalent bonds, to determine if a compound is polar the vectors need to be examined (first look at how the the compound will look based on the repulsion of electrons, then look at which element of the compound is more electronegative and draw arrows from the less electronegative elements to the greater electronegative element, add the vectors [from tip to tail] and draw a straight line from the first tail to the final to get a resultant vector, if the vectors all cancel out then the compound is not polar, otherwise it is)

-coordinate covalent: when one atom brings both of the electrons to a bond, most notably when H+ (which has no electrons) combines with other compounds (that appear complete) to create a positively charged compound (called a ________ ion), a coordinate bond is most often shown with an arrow connecting the electron donor to the electron receiver (Figure 1)

-covalent bonds can be single, double, or triple depending on how the electrons are shared, Figure 2 shows different situations that result in these different kinds of bonds. Oxygen, for example, has six valence electrons, thus by gaining two more electron oxygen would have a complete outer shell. If two oxygen atoms each share two electrons with each other this electron configuration would be achieved. So oxygen creates a double bond with two shared paired of electrons.

-allotropes: elements that can exist in two or more physical forms, for example diamonds (very hard, not conductive, carbons are arranged in a lattice structure, it is a stable network of covalent bonds and hexagonal rings, each carbon in it is connected to four other carbons, they don't generally react to strong bases or acids) , graphite, (has a layered planar structure, each carbon is generally to three other electrons, which makes graphite a good electrical conductor) and fullerene (carbon that takes the form of a sphere, 60 carbons that combine in pentagonal covalently bonded carbon rings)

Figure 1. An example of how coordinate covalent bonds look

Figure 2. Examples of how covalent bonds can result in single, double, or triple bonds

-different bonds (single, double and triple) have different strengths that follow the trend: triple > double > single (triple being the strongest bond, and single being the weakest bond)

-different types of bonds have different bond lengths that follow the trend: single > double > triple (single being the longest bond, and triple being the shortest bond)

Lewis Diagrams of Covalent Bonds:

-Lews electron dot structures are used to show covalent bonds, they include just the valence electrons of a molecule

-some structures need to be drawn with valence shell electron pair repulsion (VSEPR) taken into consideration. VSEPR is the policy that each pair of electrons will be repelled from each other as far as possible in 3D space because electrons carry a negative charge. The repulsion levels, and thus the angle drawn in a diagram, will differ depending on the relative magnitude of the repulsion. The magnitude is as follows:

lone (not bonded) pair to lone pair > bonded pair to lone pair > bonded pair to bonded pair

-this means that if a molecule has one lone pair the other bonded pairs will be pushed away (Figure 3), or if a molecule has two lone pairs the other bonded pairs will be pushed from these two and the lone pairs will consume be as far apart as possible

-Lewis dot structures are generally one dimensional, but to show the that the molecules really exist in 3D sometime other lines are used. A long triangle as a bond shows that the molecule attached is coming forward from the plane, while a triangular dashed line shows that a molecule is attached behind the plane. This is useful for showing the true covalent structure of a compound

Figure 3. An example of a Lewis diagram that considers VSEPR

Covalent Structures:

-Lewis dot structures are generally only one dimensional, but the covalent structures can be used to describe what the structure actually looks like

-Figure 4 shows the different shapes covalent bond structure that can appear in depending on the number of electron groups and number of lone pairs, usually only molecules with electron groups from 2-4 are dealt with

Figure 4. A chart showing the covalent structures that can exist

Ionic Bonding:

-intramolecular force

-metals bonding with non-metals (Figure 5)

-cations are the metals, they are the atoms that lose electrons to form positive ions

-anions are the non-metals, they are the atoms that gain electrons to form negative ions

-ionically bonded molecules have a high melting and boiling point because a large amount of energy is needed to break the strong bonds of an ionic compound

-volatility is the tendency for a substance to evaporate, and because of strong electrostatic forces of attraction the volatility of ionic compounds is usually low

-ionic compounds in a solid state have ions that occupy a fixed position (in the lattice), thus ions are not free to move and the compound doesn't conduct electricity

-ionic compounds that have been reduced to liquid by heating (in molten state) have free ions and can conduct electricity

-ionic compounds are soluble in polar solvents but not non-polar solvents

Figure 5. Examples of ionic compounds

There are many difference between ionic and covalent bonds as seen in Figure 6

Figure 6. A comparison of ionic and covalent bonds

Electronegativity:

-can be used to distinguish if a bond is ionic, pure covalent (non-polar), or polar covalent

-to determine the electronegativity of a a substance find the difference between to two elements of the compound

-HF can be used as an example. F has an electronegativity of 4.0 ad H has an electronegativity of 2.2, which means the compound has an electronegativity of 1.8. Using Figure 7 it can be determined what type of bond HF is.

Figure 7. A chart showing how the electronegativity of a compound dictates the type of bond it has created

Metallic Bonding:

-intramolecular force

-metals bonding to metals

-electrons are delocalized (not confined to a particular location but are able to move throughout a structure) (Figure 8)

-strong attractive forces in the bond

-properties include: high melting and boiling point, insolubility in water, good conductor of heat and electricity, malleable (ability to bend without breaking) and ductile (ability to be drawn out into a wire)

Figure 8. An example of how delocalized electrons are able to move through a structure

Vanderwaal Forces:

-intermolecular force

-when atoms have a bond that is constantly moving so one element of the molecule becomes temporarily negative and the other temporarily positive

-two elements like this that are next to another will have a weak attraction

-the vanderwaal forces increase down a column

-in non-metals there is a higher vanderwaal force as the materials get more solid, in metals it is the opposite

Dipole-dipole attraction:

-intermolecular force

-weak connection of molecules from the more negative end of one molecule to the more positive end of another (Figure 9)

Figure 9. Demonstration of dipole-dipole bonds compared to covalent bonds

Hydrogen Bonds:

-intermolecular force

-hydrogen of one molecule will attach to one of three electrons (nitrogen, oxygen, or fluorine because they are the most electronegative) of another molecule, creating a weak bond (though it is strongest bond of intermolecular forces), the hydrogen will attach to lone pair electrons of other molecules

-more then one hydrogen bond can form (two elements each with both oxygen and hydrogen can create a stronger bond because each oxygen will attach to the others hydrogen, making two bonds)

-Figure 10 shows what a hydrogen bond looks like and gives examples of hydrogen bonds

Figure 10. Example of how hydrogen bonding works and molecules that may combine through hydrogen bonding

Página 2

ENERGETICS (Unit 5)