Chemistry: 1.3 Bonding Part 2
This flashcard set explains electronegativity as an atom's ability to attract shared electrons in a covalent bond. It outlines periodic trends, such as increasing electronegativity across a period and decreasing down a group, along with the top five most electronegative elements.
Electronegativity
The power of an atom to attract the pair of electrons in a covalent bond
Key Terms
Electronegativity
The power of an atom to attract the pair of electrons in a covalent bond
Top 5 electronegative elements
Fluorine, oxygen, nitrogen, chlorine and bromine
what is the trend in electronegativity across a period
increases
why does electronegativity increase across a period
nuclear attraction on outer electrons decreases
what is the trend in electronegativity down a group
decreases
why does electronegativity decrease down a group
nuclear attraction on outer electron decreases
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| Term | Definition |
|---|---|
Electronegativity | The power of an atom to attract the pair of electrons in a covalent bond |
Top 5 electronegative elements | Fluorine, oxygen, nitrogen, chlorine and bromine |
what is the trend in electronegativity across a period | increases |
why does electronegativity increase across a period | nuclear attraction on outer electrons decreases |
what is the trend in electronegativity down a group | decreases |
why does electronegativity decrease down a group | nuclear attraction on outer electron decreases |
do noble gases have electronegativity values? | no because they don’t normally form covalent bonds |
what is electronegativity like in a non polar bond | same atom so same electronegativity electrons are evenly distributed |
what is electronegativity like in a polar covalent bond | significant difference in electronegativity covalent bond where electrons are unevenly distributed |
what is electronegativity like in an ionic bond | difference so large that electrons permanently go to one of atoms forming ions |
what does delta mean |
|
how do we show a negative charge on an atom involved in a polar bond | delta negative |
how do we show a positive charge on an atom involved in a polar bond | delta positive |
what do charges depend on | how electronegative the atom is- more electronegative- more power to attract electrons- negative |
what are hydrocarbons non-polar | carbon and hydrogen have similar electronegativity |
what do polar bonds mean in a simple molecule | the whole molecule has an uneven distribution of electrons |
what happens in more complex molecules | the dipoles of the polar bonds may cancel out |
what decides whether dipoles cancel out or not | whether the molecule is symmetrical or not |
what are intermolecular forces | forces of attraction between molecules |
what are the three types of intermolecular force |
|
which substances have van der Waals’ forces within them | all molecules and atoms |
which molecules contain permanent dipole-dipole forces (and VDW) | between polar molecules |
what is a hydrogen bond | special case of permanent dipole-dipole force |
where do hydrogen bonds occur (and VDW and permanent dipole-dipole) | where theres a delta plus hydrogen atom and either a nitrogen, oxygen, fluorine with a lone pair |
example of a molecule with hydrogen bonds | water |
what is special about intermolecular forces | molecules with hydrogen bonds also contain permanent dipole-dipole forces and VDW forces molecules with permanent dipole-dipole forces also contain VDW forces |
What are VDW forces caused by | The movement of electrons which unbalances the charge distribution within the molecule. This creates an instantaneous dipole across the molecule The instantaneous dipole is constantly forming and disappearing |
What does the dipole constantly forming and disappearing induct | A dipole in neighbouring molecules, resulting in weak forces of attraction between molecules |
VDW forces are present between all molecules but… | Not ions or metals |
What properties do non-polar molecules have | Relatively low boiling points Generally gases/volatile liquids at room temperature |
Why do bigger molecules have larger induced dipoles | They have more electrons |
What do larger induced dipoles in bigger molecules result in | Stronger VDW forces between molecules |
Where do permanent dipole-dipole forces occur | Between molecules which have a permanent dipole They occur in addition to VDW forces The delta positive end of one molecule is attracted to the delta negative end of a neighbouring molecules Stronger than VDW forces |
Where does hydrogen bonding occur | Between molecules which contain a hydrogen atom bonded to either F, O or N Between a delta positive hydrogen atom in one molecule and a lone pair of electrons in an N, O or F in a neighbouring molecule Occur in addition to VDW Strongest intermolecular force |
Why do substances which contain hydrogen bonds have higher boiling points than expected | Due to the strength of the hydrogen bonds (strongest intermolecular force) |
Why do substances which contain hydrogen bonds dissolve in water | They form hydrogen bonds with water |
Why does ice float on water | The density of ice is less than that of water. |
Why does water expand when it solidifies (ice) | As the temperature gets to 0 degrees the water molecules are held further apart by hydrogen bonds in an open lattice |
What is sublimation | When a material goes straight from a solid to a gas eg. Iodine and carbon dioxide |
Energy is needed to change a substance from a solid to a gas to… | Overcome forces of attraction |
When is a bond polar | When there is an uneven distribution of electrons |
When is a bond polar | When there is an uneven distribution of electrons |
When is a bond polar | When there is an uneven distribution of electrons |
when can an electric current flow | if there are charged particles which are free to move ie. delocalised electrons or mobile ions |
when can substances dissolve | if solute and solvent molecules attract one another |
what can ionic and polar substances dissolve in | polar solvents such as water |
what can non polar substances dissolve in | non polar solvents such as hexane |
what is the structure in ionic compounds | giant ionic lattice |
how are the ions arranged in the lattice | negative and positive ions alternate | . each ion surrounded by oppositely charged ions in all directions |
in soldium chloride, what is each sodium ion surrounded by | 6 chloride ions |
in sodium chloride, what is each chloride ion surrounded by | 6 sodium ions |
what is the giant ionic lattice held together by | strong ionic bonds |
why do ionic compounds have high melting points | lots of strong ionic bonds need to be broken |
do ionic compounds conduct electricity | when solid no when dissolved/molten yes |
why do ionic compounds not conduct electricity when solid | ions held fixed in the lattice |
why can ionic compounds conduct electricity when molten/dissolved | ions free to move |
why are ionic substances brittle | if enough force is applied the layers slide over eachother because like charges move next to eachother, causing repulsion and the lattice structure breaks down |
what is the structure of a metal | positive metal ions surrounded by a sea of delocalised electons positive ions fixed electrons free to move |
why do metals have high melting points | giant structure and metallic bonds are strong |
why can metals conduct electricity when solid/liquid | delocalised electons can flow through structure and carry the current |
why are metals strong | metallic bonds are strong and extend through the giant metallic lattice |
what does metallic bond strength depend on | size and charge of metal ion (smaller and higher charged ions are stronger) |
what is a malleable substance | can be hammered or pressed into shape without breaking or cracking |
what is a ductile substance able to do | be drawn into a wire |
why are metals malleable/ductile | layers of ions in giant metallic lattice can slide over each other into new positions without disrupting metallic bond |
what are the two main types of covalent substance | simple and giant/simple and macro-molecule |
describe the structure of a molecular crystal of iodine | covalent bonds between iodine atoms- weak VDW forces between i2 molecules |
why do simple covalent molecules have low melting points | all that Is needed is to overcome weak intermolecular forces (VDW,dipole-dipole, H-bonds) |
are simple covalent molecules soluble | usually insoluble in water unless they can form hydrogen bonds/react with water |
why don't simple covalent molecules conduct electricity | no charged partices which are free to move, don't contain ions/delocalised electrons |
examples of macromolecular crystals | diamond, graphite and graphene |
why do macromolecular crystals have a high melting point | strong covalent bonds between all atoms. Lots of energy needed to break these |
why does diamond not conduct electricity | electrons are localised in the covalent bonds so are not free to move- no ions present |
why does graphite conduct electricity | one electron per carbon not involved in bonding and is delocalised along the layer |
why does graphene conduct electricity | better conductor than silver! delocalised electrons carry current |
why are the macromolecular crystals insoluble in water | covalent bonds are very strong and the lattice does not break up when any solvent is added |