module 12.1- bonding Flashcards
why do elements in g18 exist as single atoms? what is this called?
monatomic (existing as single atoms) - due to their extreme stability, they rarely bond with other atoms
what is the charge of an atom?
no charge- they are neutral due to equal no. of protons and electrons
how are cations and anions formed?
- metal atoms lose valence electrons
- non-metal atoms gain electrons
describe metallic bonding, referencing the sea of delocalised electrons
- in metals, the valence electrons of atoms are loosely held
- these electrons detach from their individual atoms, becoming delocalised, meaning they are free to move throughout the entire metallic structure
- the remaining metal atoms form a lattice of positively charged ions
- the delocalised electrons create a “sea” that surrounds and moves freely around the positively charged metal ions
- since opposites attract, this electrostatic force provides multidimensional bonding between pos ions and ‘sea’ of delocalised electros
*this bonding holds the metal together and is known as metallic bonding
**metallic bonding is the force that holds metal atoms together, with freely moving electrons acting as the “glue” between the metal ions.
explain why metals are malleable, ductile & conductors of heat/electricity
- when stress is applied (such as bending or stretching), the layers of metal ions can slide past each other without breaking the metallic bond. this movement is possible because the delocalised electrons can move with the metal ions, maintaining the electrostatic attraction between them.
*this ability to allow the atoms to shift positions without breaking the overall structure makes metals malleable and ductile. - electrical conductivity: the delocalised electrons are free to move throughout the metal lattice. when an electric potential is applied (e.g., when a metal is connected to a power source), these electrons can flow through the metal, carrying charge and allowing electrical current to pass
- thermal conductivity: the delocalised electrons also help transfer energy. when one part of the metal is heated, the free electrons gain kinetic energy and move rapidly. they then transfer this energy to other parts of the metal, efficiently distributing heat.
what force pulls cations and anions together?
electrostatic attraction
can ionic substances conduct electricity?
- when solid, ionic substances cannot conduct, as the ions are bonded within their lattice= cannot move freely
- however, when molten or dissolved in water, ions become mobile as they separate from each other= allows them to conduct an electric current
when does a covalent bond occur?
when two non-metals share 1+ pairs of valence electrons
- if one pair is shared, then one electron from each atom forms the bond (holds the 2 atoms together)
what do covalent bonds usually result in?
- the formation of groups of atoms known as molecules
what is electronegative a measure of?
- a measure of how strongly an atom pulls electrons towards it
e.g. atoms of non-metals pull e- towards them more strongly than atoms of metals= non-metals are more electronegative than metals
what happens to atomic radius from left->right?
decreases
distinguish polar from non-polar bonds
- if there is a diff. in electronegativity between 2 atoms, the pair of e- will spend more time near the electronegative atom= polar bond
- when the pair of e- is distributed evenly between 2 atoms= non-polar bond
what does the symbol delta represent?
- partial charge in a polar bond
define electron cloud
- regions of negative charge that surround the nucleus of an atom *where e- are found
describe the relationship between net dipole and polarity
- if a molecule contains one polar bond= net dipole
- if a molecule contains multiple polar bonds= **net dipole OR non-polar
–> if direction of polar bonds is NOT cancelled out= net dipole
–> if direction of all polar bonds is EXACTLY OPPOSITE= non-polar
what are the 3 types of intermolecular bonding- provide an explanation of each + general info about intermolecular bonds compared to other bonds
strongest to weakest: hydrogen > dipole-dipole > dispersion
*all intermolecular bonds are far weaker than covalent, ionic or metallic bonds
-
dispersion forces *temporary dipole
-> weak strength
-> exist in all molecules
-> e- in atoms and molecules are constantly moving. e- might become unevenly distributed, creating a temporary imbalance of charge (a temporary dipole).
this can induce a similar dipole in a nearby molecule, creating a weak attraction between the two.
-> larger molecules or atoms with more electrons= stronger dispersion forces
**weak attractions caused by temporary shifts in electron distribution, allowing molecules to attract each other momentarily. responsible for the intermolecular forces in non-polar molecules -
dipole-dipole bonding *permanent dipole
-> occur between polar molecules (polar bonds and asymmetrical) but not between non-polar molecules
-> much stronger than dispersion forces as the attractive force is greater
-> the greater the polarity of a bond (the difference in electronegativity), the stronger the dipole and the stronger the dipole-dipole bond
-> the partially positive end of one molecule is attracted to the partially negative end of another molecule -
hydrogen bonding
-> a type of permanent dipole-dipole attraction that only occurs between molecules in which hydrogen is bonding directly to one of nitrogen, oxygen or **fluorine*
-> strong because of the large difference in electronegativity between H and N, O and F
-> determines melting and boiling point
(hydrogen bonds are always drawn as dashed lines)
define meniscus
- the curve at the surface of a liquid when it is in contact with a solid, such as the inside of a glass container
*liquid’s surface is slightly curved due to surface tension, so reading from the bottom of the curve gives accurate liquid level.
