bonding Flashcards
substances
made of single repeating unit at atomic level
mixture
made of several repeating units at atomic level
2 properties of particles
always have kinetic energy
always hold onto each other
3 properties of metals
malleable
ductile
conductive
electron sea model
electrons in a metal can move freely around in the structure
what keeps electrons together in electron sea model
positive ions form structure and balance charge of delocalized electrons
whats a lattice
regular repeating structure of ions
what holds lattice together
electrostatic force between cation and electron
stronger attraction in cations and electrons
stronger metallic bonding
why are metals malleable
lattice shape changed but cations can slide over each other keeping shape and metallic bonding
what happens when metal turns to liquid
regular lattice structure broken and cations can slide over each other
properties of liquid metal
not malleable or ductile
shiny and conductive
bonding in a liquid metal
does have delocalized electrons
weaker metallic bonding
higher melting point of metals
stronger metallic bonding
properties of gas
no metallic bonding and higher boiling point
do metals in liquid state have lattice
no lattice of cations
what happens when a metal boils
overall attractive forces overcome repulsive forces so no balance of forces
why do metals have higher melting and boiling points
because there is a strong attraction between cations and delocalized electrons
properties of metal gases
only have metal atoms
trend of metallic bonding
increases across period
why do metallic bonding increase across period
ions in lattice become more charges and number of electrons increases so higher electrostatic force between cation and electrons
why cant metals delocalize electrons all the time
only delocalize outer shells electrons
why can metals only delocalize outer shell electrons
only largest orbital overlap
small orbitals dont overlap so cant delocalize
trend of metallic bonding down group
decreases
why does metallic bonding decrease down group
size increases so more distance between cations nucleus and electrons so weaker forces
2 features of cations that effect metallic bonding
charge and size
state and explain general properties of metals exam question
melting point across period 3 eg
charge of cation increases so stronger attraction between cations and electrons
stronger attraction means stronger metallic bonding
compare melting point of 2 metals
eg mg ions more charges than potassium ions so stronger attraction of cations and electrons
stronger bonding so more energy to break forces
explain general melting point down group
size increases so more distance between cation and electrons so less energy required
what are alloys
metals with 2 or more elements
properties of all salts
high melting and boiling points
brittle
soluble
only conductive soluble
whats the perfect ionic model
all salts are made of positive and negative ions arranged in 3D lattice and ions perfect spheres
what keeps lattice together in ionic bonding
negative and positive electrostatic force/ attraction
what is ionic bonding
attraction between all cations and anions in lattice
charge of silver
+
charge of zinc
2+
charge of iron
2+ or 3+
charge of copper
1+ or 2+
charge of lead
2+ or 4+
group 7 6 5 anions
7= -
6=2-
5=3-
how to name anions
add ide to element name
hydrogencarbonate ion
HCO3-
sulfate ion
SO42-
phosphate ions
PO43-
maganate ions
MnO42-
hydroxide ions
OH-
how do cations and anions combine
1 to 1 ratio to keep compound neutral
how to present ionic compounds
use empirical formula
how ionic compounds are formed in terms of electrons
how many electrons transferred from one element to other
what happens when ionic compound melts
lattice breaks down and flow over each other
weaker ionic bonding
liquid ionic compound into gas
seperates into individual compounds
no ionic bonding
properties of solid ionic compounds
cant carry charge
dont conduct electricity
properties of liquid ionic compound
can carry charge
can conduct electricity
properties of gas ionic compiund
cant carry charge or conductive
why are ionic bonds so strong
attraction of negative and positive
repulsion between both same
why do ionic compounds have high melting and boiling points
strong electrostatic attraction between anions and cations
whats is the electrostatic attraction effected by
charge and distance between particles
what happens as you move down group for ionic bonding
elements get larger
ioselectronic
equal number of electrons
property of isoelectronic ions
larger nuclear charge smaller ion
how do predict melting points
look at cations and anions and compare charge and then size
why are ionic compounds brittle
when hit opposite charges together and lots of repulsion breaks layers apart
how does a covalent bond form between two atoms
shared pare of electrons between atoms
how much covalent bonds to atoms form
until they have full outer shell
problems with circle and dot diagrams
cant see which electron came from where
messy
bonding pairs
pairs of electrons bonding in covalent bonds
lone pairs
outside unbonded electrons
double bonds
2 pairs of electrons shared between 2 atoms
triple bonds
3 electrons shared between 2 atoms
how to predict which covalent bond will be longer
compare size of atoms forming bonds
bigger atoms is longer bonds
how to predict melting and boiling point of bonds
compare bond lengths
shorter lengths stronger
dative covalent bonds
one atom contributes both electrons to the shared pair
how to represent dative bond
arrow from which one have it to where its going or two of same dot or cross
2 elements to form electrons
dative bond donors and acceptors
valence bond theory
if two atoms have orbitals which have unpaired electrons then orbitals will overlap with each other forming a covalent bond
how do orbitals overlap
one pair of orbitals overlap head on to form one bond and other bond side on
sigma bonds
when orbitals overlap head on lower case sigma
pi bonds
orbitals overlap side on
why does carbon form 4 bonds instead of 2
ground state = lowest energy orbital
in ground state can have 4 unpaired electrons
electronegativity
ability to attract electron in covalent bond
electronegativity across period
increases as more protons
electronegativity down group
decreases down group
shell increases so nucleur radius bigger so less atraction
electronegativity
ability to attract electron pair in a covalent bond
factors that affect electronegativity of bonded pair
-nuclear charge
-distance from nucleus
-shielding
exceptions of trends of electronegativity
-hydrogen similar to C as no electron shielding
-dont apply to group 8
-rare bonds
dipole
one part of molecule slightly negative and the other slightly positive
polar
slightly negative and positive parts of molecule
comparing polarity
look at pauling scale values and compare difference
ionic compounds formed when electronegativity high
when difference between to atoms so high the electrons almost belong to that atom
-atom able to conduct electricity and metal and brittle
-electrons sent to other atom
metals formed in electronegavitity
when difference lower than 1.5 or greater
what holds polyatomic ionic bonds together
covalent bonds
ionic compounds with polyatomic ions
covalent and ionic bonds
accounting for charge when drawing polyatomic ions
if compound has negative charge add electrons to most electronegative compound
-if molecule positive add electron to least electronegative molecule
