Chem 1 - Bonding

Metal and Non-metal

Transfer of electrons from metal atom to non-metal atom producing charged ions.

• metal ---> +ve ions (electrons lost)

• non-metal ---> -ve ions (electrons gained)

Electrostatic Attraction - between the +ve and -ve ions = ionic bond

Properties:

• High melting/boiling point - due to structure and electostatic attraction

• Conducts electricity if dissolved in water/molton

Q1.) In what form does an ionic compound conduct electricity?

Q2.) Explain in terms of electrons the structure of an ionic bond?

Q3.) Why do ionic compounds have a high boiling point?

Q4.) Draw the structure of sodium chloride (including charges/atom sizes)

Q5.) Draw a dot and cross diagram of magnesium oxide formed by a magnesium atom and an oxygen atom?

Q6.) Explain why magnesiums electrostatic force of attraction is stronger than sodium? (hint: charge)

Non-metal and Non-metal

Sharing electrons to obtain a full outer shell. (since there is no transfer of electrons there is no charge particles - ions)

Covalent bonds are very strong

Dative Covalent (or co-ordinate) same strength

e.g.   NH3 + H+

The lone pair of electrons on the NH3 is attracted to the H ion, the covalent bond is represented by a arrowed. (draw in displayed formula)

  NH3 ---> H    =  [ NH3 --->]+

Q1.) Explain in terms of electrons what is meant by a covalent bond?

Q2.) Why do atoms want to form covalent bonds?

Q3.) Give an example of a covalenty bonded compound?

Q4.) Draw a dot and cross diagram for the formation of methane?

Q5.) Draw a dot and cross diagram for the dative bond between NH4?

Q6.) Explain why the NH3 molecule is attracted to the H+ ion?

Q7.) How strong is a dative covalent bond?

Electronegativity - a measure of tendency an atom has to attract a bonding pair of electrons in a covalent bond.

Pauling Scale - compares electronegativity. Most electronegitive elements: fluorine, oxgyen, nitrogen (FON)

Polar Bond - a covalent bond in which there is a seperation of charge, making one end delta + and other delta -  (e.g. H2O, HCl)

                          (delta +)  H          :  Cl (delta -)

Chlorine is more electronegative then hydrogen so attracts bonding pair of electrons closer to itself.

Q1.) Define the term electronegativity?

Q2.) What are the most electronegative elements?

Q3.) Define what is meant by a polar bond?

Q4.) Explain why/how one element becomes delta - and one delta +?

Q5.) If the bond is pure covalent this means theres is no polarity in the bond draw an example for this?

Q6.) What would this mean the electronegativity of the elements are in terms of comparison?

Molecular: simple molecules (NH3, H2O)Bonding capcity reached within few atoms.

Structure of Iodine - v.regular structure, mixture of covalent bonds & intermolecular forces

Properties

low boiling/melting point - weak intermolecular forces need less energy to overcome.

No electricial conductivity - no delocalised electrons.

Most substances are soft, crumble not very strong.

Q1.) Explain why iodine has a low boiling point?

Q2.) Draw the structure of iodine?

Q3.) Define molecular? give an example of a molecular compound?

Q4.) Can molecular compounds conduct electricity? explain?

Q5.) If iodine is a gas explain what would happen to its structure in terms of intermolecular forces and covalent bonds?

Forms a lattice (carbon, silicon, silicon dioxide)

Diamond

• Carbon - 4 bonds each     v.hard, brittle, strong - covalent bonds
  
• v.regular structure             are strong and directional
  
• 3D                                   Uses - drills, glass cutting
  

Properties:

v.high melting/boiling point - lots of energy needed to break the many strong covalent bonds the melting/boiling point are v.simular - directional (once broken completely broken)

Electrical Conductivity - none - no delocalised electrons

Q1.) What is a giant covalent structure?

Q2.) Draw the structure of diamond?

Q3.) Why does diamond have a very high melting point?

Q4.) Why is the melting and boiling point of diamond very simular?

Q5.) State and explain the properties of diamond?

Q6.) Name some uses for diamond? explain why it is used for this?

Q7.) Why is diamond brittle?

Graphite - forms a lattice

• 2D                                                uses: pencils, lubricant
  
• Layers held together by intermolecular forces
• Layered (hexgons) each carbon atom is bonded to 3 other crabon atoms, the spare electron is delocalised and stays in the space between layers.

Properties

Electrical conductivity - Yes, due to delocalised electrons in each plane Density - lower than diamond - due to larger distances between planes Hardness -much softer than diamond - the planes can slip over each other

Q1.) Explain in terms of structure why graphite is used as a lubricant?

Q2.) Draw the structure of graphite

Q3.) Explain the structure of graphite in terms of bonding?

Q4.) Can graphite conduct electricity? explain?

Q5.) Explain what is meant by a delocialised electron?

Q6.) What are the uses of graphite? explain why they are used for this?

Q7.) In terms of bonding explain the properties of graphite?

Metals only - atoms are packed closely together in 2 possible ways

• Atoms in gaps
• Atoms directly under - no gaps

The electrons feel an attraction to the nucleui of other atoms surrounding and become delocalised, once it leaves the original atom a +ve ion is formed. Metallic Bond - an array of +ve ions and delocalised electrons

3 examples -              Na          --->         Na+    +    e-

                                Mg          --->         Mg2+  +   2e-

                                Al            --->         Al3+   +   3e-

Higher charge = higher melting point - stronger attraction between +ve ions and delocalised electrons.

e.g. Mg has double delocalised e- per +ve ion compared to Na, so more energy needed to break attraction - higher boiling/melting point.

Properties

Good conductors of heat - closely packed particles vibrate passing on energy

Good Conductors of electricity - delocalised electrons move carrying the current (electricity = moving electrons)

Strength - depends on size of atoms and charge

More Properties

Metals are malleable - can be bent/shaped

Metals are ductile - can be stretched - wires

High melting/boiling point - strong attraction between +ve ions and delocalised electrons.

Questions

Q1.) Define metallic bond?

Q2.) Draw the 2 ways atoms in a metal can be arranged?

Q3.) Explain how a metallic bond is formed?

Q4.) Explain why Aluminum has a higer boiling point than sodium?

Q5.) Explain why metals are good conductors of electricity?

Q6.) Explain why metals are good conductors of heat?

Q7.) What does ductile mean?

Q8.) What effects the strength of a metal?

Q9.) Why do most metals have a high melting point?

Q10.) Explain in terms of structure and bonding the properties of metals?

Q11.) What does malleable mean?

Q12.) Give a use for a ductile metal? explain?

• Van Der Waals' forces (VDW)
• Dipole - Dipole                                 3 types
  
• Hydrogen Bonds

Van der waals' forces - exists in all simple molecules/atoms - these electron are constantly moving, so distribution around the nucleus is unlikely to be even.

cloud of electrons         VDW attraction

                           delta +   ------------- delta -

uneven distribution = temporary dipole constanty changing - fluctuating dipoles. but overall force remains. Size of atom determines strength of VDWs' (more electrons = stronger VDWs' forces)

Dipole - Dipole

Molecules with permant dipoles e.g. HCl (H = delta +, Cl = delta -)

              H - Cl --- H - Cl ---H                --- dipole attractions

Chlorine is more electronegative then hydrogen

Hydrogen Bond (strongest intermolecular force)

• Strength 1/10 of a covalent bond
• occurs when hydrogen is bonded to a FON element

HF, H2O - 2 lone pairs, 1 lone pair attracted to the H forms hydrogen bond. The 2 bonding pairs are held closely to the oxygen - due to electronegativity. The single proton in H nucleus attracted to +ve ion.

Q1.) Where are van der waals' forces found?

Q2.) List the 3 types of intermolecular forces by weakest - strongest?

Q3.) What is an example of a molecule containing dipole-dipole forces?

Q4.) Draw the attraction between hydrogen atoms and Chlorine atoms in HCl?

Q5.) What is the strength of a hydrogen bond in camparison to a covalent bond?

Q6.) Draw a cloud of electrons and temporary dipoles to represent van der waals' forces?

Q7.) What effects the strength of a van der waals' force?

Q8.) Draw how hydrogen bonds are formed in 2 H2O molecules?

Q9.) What is attracted to the hydrogen atom to form a hydrogen bond?

Q10.) What intermolecular force is present in HF? draw a diagram to represent this?

Q11.) Why do dipole - dipole forces form?

Q12.) HCl has a lower boiling point compared to HF explain why?

Q13.) In terms of intermolecular forces explain why group 7 elements boiling points increase down the group?

14.) Why do the distribution of electrons create temporary dipoles?

Bonding - covalent - total number of electron pairs dictate molecular shape.

2 pairs = e.g. BeCl2         Linear                      Bond Angle = 180

3 pairs = e.g. BF3           Trigonal Planar          Bond Angle = 120

4 pairs = e.g. CH4           Tetrahedral               Bond Angle = 109.5

5 pairs = e.g. PF5           Trigonal Bipyramid    Bond Angle = 120

                                                                     180 in other plane

6 pairs = e.g. SF6           Octahedral                Bond Angle = 90

Q1.) Why is each angle exactly 120 in BCl3?

(ANS- equal repulsion between the bonding pairs of electrons in all 3 covalent bonds)

Q2.) Predict the bond angle in CCl2? (ANS- 120 - 2 (2 lone pairs of e-) = 118)

Q3.) In terms of electron pairs explain why the bond angle in NH4+ ion are all 109? (ANS- equal repulsion between bonding pairs of e-'s from all 4 covalent bonds around nitrogen atom in a tetrahedral shape)

Q4.) From 2 pairs to 6 pairs of electrons name the shape, give an example and the bond angle/s.

Q1.) Explain why HF is polar covalent?

(F is more electronegative than H ... bonding e-'s attracted to F)

Q2.) What is meant by polarised?

(Electron cloud is unevenly distributed)

Q3.) State a feature which would cause an ion to be polarised strongly?

(Higher charge)

Q4.) Explain why the melting point of silicon is very high?

(giant covalent...many strong covalent bonds...need lots of energy to overcome)

Q5.) Explain why the shape of the NH4+ ion is a regular tetrahedral?

(equal repulsion between bonding pairs in all 4 covalent bonds)

Q6.) Describe the bonding in a metal? (metallic bonding...an array of positively charged ions...attracted to delocalised electrons...closely packed)

Q7.) Explain why magnesium has a higher melting point than sodium?

(Mg has a higer charge Mg+2 ions meaning there are double delocalised electrons per +ve compared to Na)

Q8.) Sodium chloride can be form by reacting sodium with chlorine, write an equation for this reaction?

(Na + Cl2 ----> 2NaCl)

Q9.) A chloride ion has one more electron than a chlorine atom, in the formation of sodium chloride, where does this electron come from? (the sodium atom)

Q10.) Define electronegativeity?