(Chem 2) Quiz 1: Chapter 9 (Chemical Bonding: The Lewis Model)
What is *formal charge*?
During bonding, atoms may end with more or fewer electrons than the valence electrons they brought in order to fulfill octets. This results in atoms having a formal charge.
What happens in a covalent bond?
Electrons are *shared* between nonmetals - nonmetals have high ionization energies, so hard to form ionic bonds (requires a lot of energy) - so better to form covalent bond because can lower its potential energy by getting the electrons in between the nucleis of the two atoms - redistributes electron density to get it further away from the nuclei of its respective atom
What happens in covalent bonding?
Electrons are *shared* between nonmetals - nonmetals have high ionization energies, so hard to form ionic bonds (requires a lot of energy) - so better to form covalent bond because can lower its potential energy by getting the electrons in between the nucleis of the two atoms - redistributes electron density to get it further away from the nuclei of its respective atom
What happens in an ionic bond?
Electrons are *transferred* between metal and nonmetal
What are *lone pairs*?
Electrons that are not shared by atoms but belong to a particular atom are called lone pairs (also known as nonbonding pairs)
What are *bonding pairs*?
Electrons that are shared by atoms are called bonding pairs
Which atoms violate the octet rule when forming compounds?
- Be generally likes to form two bonds and no lone pairs in its compounds - B generally likes to form three bonds and no lone pairs in its compounds - many elements (3rd row and below) may end up with more than 8 valence electrons in their structure because they can use their empty d orbitals to place excess electrons
What are some general rules for drawing Lewis structures?
- Carbon always wants to make 4 bonds and have 0 lone pairs - Nitrogen wants to make 3 bonds and have 1 lone pair - Oxygen wants to make 2 bonds and 2 lone pairs - Hydrogen and halogen want to make 1 bond Often, Lewis structures with line bonds have the lone pairs omitted (i.e H20)
What are the exceptions to the Octet rule?
- H, Li, Be, B try to get electron configuration like He (He only has 2 valence electrons, a *duet*, but is stable because has filled 1s-shell) - sometimes Be will covalently bond, resulting in 4 valence electrons - B loses 3 electrons in covalent bonds, resulting in 6 valence electrons Expanded octets for elements in period 3 or below (may get more than 8 valence electrons. ie, compound PCl5 - phosphorus with 5 covalent bonds = 10 valence electrons. this is possible because elements below period 3 have room for d-orbitals) - results in filling up s-orbital and p-orbitals of whatever their outermost energy level is
What does the Lewis theory allow us to predict?
- allows us to predict the distribution of valence electrons in a molecule (useful for understanding the bonding in many compounds) - allows us to predict shapes of molecules - allows us to predict properties of molecules and how they will interact together
What are trends in bond energies?
- as increase the number of covalent bonds between the atoms, it becomes harder to break the bonds and releases more energy when you form them - the smaller the atomic radius, the shorter the covalent bond = the stronger the bond
What are the exceptions to the Octet rule?
- elements with empty d-orbitals can exceed the octet rule (3rd row down) - for those that can't fill an empty d-orbital, i.e. NO that has total of 11 valence electrons, will get an odd number. 1 has to be unpaired = *free-radical*, very reactive & unstable (antioxidants bind with free radicals to prevent it from being dangerous to body)
What are the trends of electronegativity?
- increases going across a row from left to right - decreases going down a column - flourine is the single most electronegative element - does not apply to noble gases (already stable so not interested in drawing in electrons to itself) - the larger the difference in electronegativity, the more polar the bond
What interactions need to be considered to calculate the potential energy when a bond forms?
- nucleus-to-nucleus repulsions - electron-to-electron repulsions - nucleus-to-electron attractions
What are the properties of molecular compounds (covalent bonding)?
1. *Hydrogen, nitrogen, oxygen and the 4 halogens (fluorine, chlorine, bromine, iodine) are all diatomic molecular elements* (as predicted by Lewis theory) - Oxygen is generally looking for two single bonds or a double bond with other atoms (i.e. water) 2. *Compounds of nonmetals are made of individual molecule units (vs. ionic compounds that have a long-range repeating lattice structure)* (i.e. water is composed of individual H20 molecules - it is not a network) 3. *Molecular compounds have low melting points and boiling points (compared to ionic compounds).* - i.e. easier to boil H20 because not breaking apart the molecules, but are breaking apart the bonds attracting the molecules together vs. in ionic compounds - if have block of NaCl - want to vaporize - need to break apart the bonds holding the individual bonds together = requires a lot more energy 4. *Molecular compounds are found in all 3 states at room temperature (vs. ionic compounds are only solids at room temp) 5. *Some molecular solids are brittle and hard, but many are soft and waxy (more variety compared to ionic that are always very hard and brittle due to fixed structure)* - i.e., diamond = carbon, a molecular compound 6. *Molecular compounds do not conduct electricity when in its solid or liquid state* (as Lewis predicted, because there are no transfer of electrons- so there are no ions in molecular compounds, no charged particles moving around) - EXCEPTION: molecular acids conduct electricity when they are dissolved in WATER, but not in the solid or liquid state, due to them being ionized by the water - give up their H+ ions) - if have aqueous solution of HCl - do not stick hand in it (but if HCl in powder/solid state - will be non-harmful) 7. *The more electrons two atoms share
What are the properties of ionic compounds?
1. *Ionic compounds have high melting points (above 300C) and boiling points (hard to break the attraction)* 2. *all ionic compounds are solids at room temperature* - Lewis theory implies that the attractions between ions are strong (and therefore, hard to break apart the attraction) 3. *Ionic solids are relatively hard* (i.e. salt crystals) - Lewis theory predicts that moving ions out of place should be difficult because completely surrounded, and ionic solids should be hard 4. *Ionic solids are brittle.* When struck they shatter. (they are hard but also very brittle) - Lewis theory predicts ionic solids will be brittle if the ions are displaced from their position in the crystal lattice (repulsive forces will occur) 5. *Ionic solids do not conduct electricity.* - Lewis theory predicts that ionic solids should not conduct electricity (because ions are held together in place pretty strongly in the crystal lattice structure - not flowing freely) (i.e. salt) - To conduct electricity, a material must have charged particles that are able to flow through the material 6. *Ionic compounds conduct electricity in the liquid state or when dissolved in water.* - the ions are separated and freely moving (i.e. add water to salt) - Lewis theory predicts that both a liquid ionic compound and an ionic compound dissolved in water should conduct electricity.
What are the properties of metallic bonding?
1. *metallic solids conduct electricity well* - theory implies that because the electrons are delocalized in sea, they are able to move freely through the metallic crystal 2. *as temperature increases, the electrical conductivity of metals decreases* - as temperature increases, atoms start vibrating faster, makes it harder for electrons to travel through the crystal - i.e. facilities with super computers have to keep a very cool temperature, because conductivity will go down and speed of computers will slow (immersed in cool non-conductive fluids) 3. *metallic solids conduct heat well* - have enormous pool of delocalized electrons and because electrons are very small, its easy to increase their kinetic energy and get them moving fast. Thus, increases temperature of the metal - i.e. cookware made of metal 4. *metallic solid reflect light* - emit light when jump from higher energy level to lower energy level - large sea of delocalized electrons, so very prone to absorbing photons of light and re-emitting photons of light when excited electron jumps back down 5. *metallic solids are malleable and ductile* - if shift position of metals by hammering, the sea of electrons will go with it. it's flexible, adaptable to changes in structure 6. *metals generally have high melting points and boiling points* - because there are strong attractions holding the metallic crystal together and these must be broken 7. *melting points of metal generally increase left to right across a row* 8. *Melting points of metals generally decrease going down a column*
How many valence electrons do the transition metals have?
All transition elements have either one or two valence electrons - except for the 5 irregulars
What is the *Octet rule*?
Atoms will bond with other atoms to get 8 valence electrons
How is bond strength measured?
Bond strength is measured by how much energy must be added into the bond to break it in half - the more electrons two atoms share, the stronger the bond should be - triple bonds are stronger than double bonds, and double bonds are stronger than single bonds (i.e. harder to break Nitrogen gas compared to oxygen gas and hydrogen gas because it has a triple bond. Oxygen has a double bond, and hydrogen has a single bond)
Why do atoms bond?
Chemical bonds form because they lower the potential energy between the charged particles that compose atoms. - elements and compounds want to lower internal energy they have stored because not stable to have a lot of stored energy - wants to release energy to be in more stable state - more stable = lower potential energy
What are *bond energies*?
Chemical reactions involve the breaking of bonds in reactant molecules and making new bonds in the products - if know how much energy it takes to break a single bond and know how much energy is released to form a bond, can add them together to get total energy of the reaction
What is a *polar covalent bond*?
Covalent bonding between nonmetals that have different electronegativities - results in unequal sharing of electrons - the more electronegative one will have higher electron density
What is the formula for calculating formal charge?
FC = valence electrons in ground state - # of lone pair electrons - # of covalent bonds REVIEW PRACTICE PROBLEMS
Where will electrons move in the formation of an ionic bond?
From the metal to the nonmetal. The nonmetal is going to steal electrons from the metal to form an anion, the metal will form a cation, and then they will be electrostatically attracted to each other (opposite charges attract)
What is the *percent ionic character*?
How closely a bond resembles an ionic bond The percentage of a bond's measured dipole moment compared to what it would be if the electrons were completely transferred - the higher the %, the more it closely resembles an ionic bond - the % will increase as the difference between the 2 atoms electronegativities reaches the 2.0 barrier - given covalent bond and differences, which bond will have greater % ionic character?
If the difference in electronegativity between bonded atoms is *0.1-0.4*, what type of bond is formed?
If the difference in electronegativity between bonded atoms is *0.1-0.4*, the bond is *nonpolar covalent* (there will be a certain amount of unequal sharing but not enough to turn it into a genuinely polar bond)
If the difference in electronegativity between bonded atoms is *0.5-1.9*, what type of bond is formed?
If the difference in electronegativity between bonded atoms is *0.5-1.9*, the bond is *polar covalent* (the degree of unequal sharing is high enough)
If the difference in electronegativity between bonded atoms is *larger than or equal to 2.0*, what type of bond is formed?
If the difference in electronegativity between bonded atoms is *larger than or equal to 2.0*, the bond is *100% ionic* (no more sharing as the more electronegative one will steal more electrons - electrons were transferred)
If the difference in electronegativity between bonded atoms is 0, what type of bond is formed?
If the difference in electronegativity between bonded atoms is 0, the bond is *pure covalent* (when atoms are identical or have identical electronegativities)
What are the 3 types of bonds?
Ionic Covalent Metallic
What does lattice energy depend on? *NEED TO KNOW*
Lattice energy depends directly on the size of the charges of the ions and inversely on distance between ions
What is the Lewis Bonding Theory?
Lewis Structure - using dots around the atom symbol to represent the valence electrons - pair the first two dots to represent the electron in the s-orbital - put one dot on each open side for the first 3 p electrons - then, pair the rest of the dots for the remaining p electrons Atoms bond because it results in a more stable electron configuration. - more stable = lower potential energy
How is strength of bond related to length of bond?
Lewis theory predicts that the more electrons two atoms share, the shorter the bond should be (= the stronger the bond) - triple bond is the shortest STRENGTH OF BOND IS DIRECTLY RELATED TO ITS LENGTH
What happens in a metallic bond?
Metals bonded to other metals (*delocalization of electrons*) - metals tend to have lower IE (easier to give up electrons) - metal atoms releasing its valence electrons to be shared as a pool by all the atoms in the metal - series of ion arranged in 3-D lattice and a sea of *delocalized* electrons spread across entire surface of the metal (which is why metals are good conductors - easy for additional electrons to float in and out of that sea) - electrons aren't localized around an individual atom
What is the stable arrangements of metals and nonmetals?
Metals form cations (lose electrons) to form the stable electron configuration of the nearest noble gas. Nonmetals form anions (gain electrons) to form the stable electron configuration of the nearest noble gas.
What is the 4 types of bonds based on the difference in electronegativity? *NEED TO KNOW*
Pure covalent Nonpolar covalent Polar covalent Ionic
How can enthalpy change of a reaction from bond energies be estimated?
Step 1: Add energies need to break each bond Step 2: Add energies released to form each new bond Step 3: Subtract. Total energy to form new bonds - Total energy to break bonds Net energy change (positive) that was released during course of reaction = exothermic (energy released) REVIEW EXAMPLE
What does the *crystal lattice* do?
The crystal lattice maximizes the attractions between cations and anions, leading to the most stable arrangement - negative charges completely surrounded by positive charges and positive charges completely surrounded by negative charges = stable
Where does the extra energy that is released in an ionic bond come from?
The extra energy that is released comes from the formation of a structure (the *crystal lattice*) in which every cation is surrounded by anions, and vice versa. Thus, ionic bonding is an exothermic process.
What happens in an ionic bond formation?
The ionization energy of the metal is endothermic - stealing an electron away from the metal to form a cation - forming the cation is an endothermic process (because need to put energy in to steal the electron away) The electron affinity of the nonmetal is exothermic - nonmetal is accepting an electron to form an anion - forming an anions is an exothermic process Forming the cation absorbs more energy than forming the anion releases. So, collectively, the formation of ionic bonds should be an endothermic process. *HOWEVER, THE FORMATION OF IONIC COMPOUNDS IS EXOTHERMIC (NEGATIVE VALUE) AND LARGE BECAUSE THERE IS EXTRA ENERGY THAT IS RELEASED*
What is the relationship between lattice energy and ion charge?
The larger charge means the ions are more strongly attracted - larger charge = stronger attraction - stronger attraction = larger lattice energy
Why are valence electrons used in bonding?
Valence electrons are the ones furthest away from the nucleus so will experience smallest effective nuclear charge - held most loosely of all the electrons in an atom - because chemical bonding involves the transfer or sharing of electrons between atoms, valence electrons are the most important electrons in bonding
What is a single covalent bond?
When atoms share one pair of electrons, it is called a single covalent bond - can be represented as a single line (each line is 2 electrons)
What is a *dipole moment*?
When there is an unequal sharing of electrons due to the nonmetals having different electronegativities (when there is a polar covalent bond) - the more electronegative one will attract more electrons to itself and gets a *partial negative charge* - not true negative charge because hasn't completely taken away electron - the end that is electron deficient (less electronegative one) gets a *partial positive charge* - not true positive charge because hasn't completely given up electron
What is a triple covalent bond?
When two atoms share three pairs of electrons the result is called a triple covalent bond - can be represented as 3 lines (total of 6 electrons, each line is 2 electrons)
What is a double covalent bond?
When two atoms share two pairs of electrons the result is called a double covalent bond - can be represented as 2 lines (total of 4 electrons, teach line is 2 electrons)
What is a dipole moment?
a measure of bond polarity (how unequally the electrons are shared) - the more electronegative atom gets the partial negative end - the less electronegative atom gets the partial positive end - depends on the difference in electronegativity between the atoms and the distance between the atoms (the longer the distance between atoms, the greater the dipole moment)
How do resonance structures exist?
as *hybrids* (like breeding dogs) - doesn't oscillate back and forth, but exists as a hybrid
What are *resonance structures*?
for Lewis structures in which it doesn't matter which side of the atom the double bond is drawn (when there is more than one Lewis structure for a molecule that differ only in the positive of the electrons) - there are 2 different Lewis structures that are equally valid - *resonance makes molecules stable* (can spread electron density out over a greater area. can delocalize that electron density so there's less repulsion) - if comparing structures - one with a resonance structure and one that doesn't - the one with a resonance structure is more stable
What is electronegativity?
how strongly an atom will pull electrons towards itself
What is *lattice energy*?
the extra energy that is released (the extra stability that accompanies the formation of the crystal lattice) - always exothermic
What is the relationship between lattice energy and ion size?
the shorter the distance between the nucleis of the ions, the more lattice energy is released = the more stable the compound is (find the one with the smallest ionic radius)
What is the *Lewis theory*?
theory used to describe how atoms interact with each other (pioneered by chemist, Lewis) - One of the simplest bonding theories - Lewis theory emphasizes valence electrons to explain bonding (valence electrons being shared) - allows us to predict many properties (molecular stability, shape, size, polarity)
