Chem Flashcards
Explain why MgO has a higher melting point than NaCl.
Mg²⁺ and O²⁻ have higher charges compared to Na⁺ and Cl⁻.
Mg²⁺ and O²⁻ are also smaller ions, allowing stronger electrostatic attraction.
This results in a stronger ionic bond in MgO, requiring more energy to break, leading to a higher melting point
Why is Al³⁺ smaller than N³⁻, even though they are isoelectronic?
Both Al³⁺ and N³⁻ have 10 electrons, but Al³⁺ has more protons (13) than N³⁻ (7).
The greater nuclear charge in Al³⁺ pulls the electrons closer to the nucleus, making it smaller.
Why do ionic compounds conduct electricity when molten but not when solid?
In the solid state, ions are fixed in place and cannot move.
In the molten state, the ionic lattice breaks down, and ions are free to move and carry charge, allowing the compound to conduct electricity.
Describe how the migration of ions can be demonstrated in electrolysis.
In electrolysis, positive ions (cations) migrate towards the negative electrode (cathode), and negative ions (anions) migrate towards the positive electrode (anode).
This movement provides evidence for the existence of ions.
As only ions can move this way in an electrical field
Why does ionic radius increase as you go down Group 1?
As you go down the group, each element has more electron shells.
This increases the distance between the nucleus and the outermost electron, making the ion larger.
What two factors determine the strength of an ionic bond?
Ionic charge: Higher charges lead to stronger electrostatic attraction.
Ionic radius: Smaller ions allow stronger attraction because electrons are closer to the nucleus
Why are ionic compounds generally soluble in water?
Water is a polar solvent with partial positive and negative charges.
The positive end of water molecules attracts the negative ions, and the negative end attracts the positive ions, pulling the ions apart and dissolving the compound.
Arrange the following isoelectronic ions in order of increasing size: N³⁻, O²⁻, F⁻.
The order is F⁻ < O²⁻ < N³⁻.
Although they all have the same number of electrons (10), N³⁻ has the fewest protons and therefore the weakest attraction between the nucleus and electrons, making it the largest.
What is meant by molar volume of gas
Space occupied by one mole of a gas at a specific pressure and temperature
What is metallic bonding?
Metallic bonding is the electrostatic attraction between positively charged metal ions (cations) and a “sea” of delocalized electrons that can move freely throughout the metal structure.
Describe the structure of metallic bonds
In metals, cations are arranged in a regular lattice surrounded by delocalized electrons. This arrangement allows the cations to stay in fixed positions while electrons move freely, contributing to the metal’s strength and conductivity.
What factors affect the strength of metallic bonds?
Charge on the metal ions (greater charge means stronger bonds).
Metallic radius (smaller radius leads to stronger bonds).
Structure of the metallic lattice (closer-packed structures have stronger bonds).
How does metallic bond strength influence melting points and density?
Metals with stronger metallic bonds have higher melting points and greater densities due to tighter packing of ions and more delocalized electrons.
Why do Group 1 metals have lower densities and melting points compared to Group 2 metals?
Group 1 metals lose one electron (ns¹), resulting in a lower charge and weaker metallic bonds. Group 2 metals lose two electrons (ns²), creating smaller ions with higher charges, leading to stronger bonds, higher density, and higher melting points.
How does the melting temperature of metals change across a period and down a group, and why?
Across a period, melting temperature increases due to smaller radii and more delocalized electrons, resulting in stronger bonds. Down a group, melting temperature decreases as larger radii and electron shielding reduce the attraction between ions and delocalized electrons.
What makes metals malleable, and how does their structure support this property?
Metals are malleable because layers of metal ions can slide over each other without breaking bonds, as delocalized electrons prevent strong repulsive forces between the ions in adjacent layers.
Why do metals have low electronegativity
Metals are more likely to loose electrons to form positive ions ,so are unlikely to attract electrons
What two factors determine the strength of ionic bonds?
Charge on the ions: Higher charges increase the attraction between ions.
Ionic radius: Smaller ions have a stronger attraction due to closer proximity, leading to stronger ionic bonds.
How does ionic radius change down a group and across a period?
Down a Group: Ionic radius increases as additional electron shells are added.
Across a Period (Isoelectronic Series): For ions with the same electron configuration, radius decreases with increasing nuclear charge, as greater attraction pulls electrons closer.
Explain the trend in ionic radius for the isoelectronic series from
N
3
−
N
3−
to
A
l
3
+
In the series
N
3
−
,
O
2
−
,
F
−
,
N
a
+
,
M
g
2
+
,
A
l
3
+
N
3−
,O
2−
,F
−
,Na
+
,Mg
2+
,Al
3+
, all ions have the same electron configuration (2,8). As nuclear charge increases from
N
3
−
N
3−
to
A
l
3
+
Al
3+
, the radius decreases due to greater attraction pulling electrons closer to the nucleus.
Why does
M
g
O
MgO have a higher melting point than
N
a
C
l
NaCl?
O
MgO has a higher melting point because
M
g
2
+
Mg
2+
and
O
2
−
O
2−
ions have higher charges compared to
N
a
+
Na
+
and
C
l
−
Cl
−
, leading to stronger electrostatic attractions. Additionally,
M
g
2
+
Mg
2+
and
O
2
−
O
2−
ions are smaller, increasing the attraction.
How do charge and ionic radius affect bond strength in ionic compounds?
Smaller ions with higher charges can pack closer together, increasing electrostatic attraction, leading to stronger ionic bonds. For example,
M
g
O
MgO has stronger bonds than
N
a
C
l
NaCl due to smaller ion size and higher charge.
Why do ionic compounds conduct electricity only when molten or dissolved in water?
In solid form, ions in ionic compounds are fixed in place by strong bonds and cannot move. When molten or dissolved, the bonds are broken, allowing ions to move freely and conduct electricity.
Describe the experiment using copper chromate to demonstrate ion migration.
Copper chromate (
C
u
C
r
O
4
CuCrO
4
) dissolved in dilute
H
C
l
HCl is layered in a U-tube with electrodes connected. When an electric current is applied,
C
u
2
+
Cu
2+
ions migrate to the cathode (turning the solution blue-green), and
C
r
O
4
2
−
CrO
4
2−
ions move to the anode (turning the solution yellow), proving the ionic nature as only ions migrate in an electrical feild
How does water’s polarity affect the solubility of ionic compounds?
Water’s polarity, with partially positive hydrogen and partially negative oxygen, attracts and separates ionic compounds. This interaction breaks the ionic bonds, causing ions to disperse and dissolve in water.
What are electron density maps, and how do they indicate bonding type?
Electron density maps show regions in a molecule where electrons are most likely found. In ionic compounds like
N
a
C
l
NaCl, contour lines don’t overlap, indicating no shared electron density between
N
a
+
Na
+
and
C
l
−
Cl
−
, which confirms ionic bonding,rather than covelant bonding
What is a covalent bond, and between which types of elements does it typically form?
A covalent bond is a strong electrostatic attraction between the nuclei of two atoms and a shared pair of electrons. It typically forms between non-metal atoms with an electronegativity of less than about 1.5 units.
What is a dative covalent (coordinate) bond, and how does it differ from a standard covalent bond?
A dative covalent bond forms when one atom provides both electrons for a shared pair. It behaves like a regular covalent bond but requires an electron-deficient atom to accept the lone pair from the donor atom.
Describe the Octet Rule in covalent bonding
The Octet Rule states that atoms bond covalently to achieve the electron configuration of a noble gas, typically eight electrons in their outer shell, leading to more stability.
Explain the concept of orbital overlap in covalent bonding, and name the types of overlaps.
Covalent bonding involves the overlap of atomic orbitals each containing a single unpaired electron. Types of overlaps include:
s-s overlap (e.g.,
H
2
H
2
)
s-p overlap (e.g.,
H
C
l
HCl)
p-p overlap (e.g., between two
p
p-orbitals in different atoms).
What is a sigma (σ) bond, and why is it strong?
A sigma (σ) bond is formed by “head-on” overlap of orbitals along the line between the centers of bonded atoms. It is the strongest type of covalent bond because the electron density is concentrated directly between the nuclei.
Why does carbon form four covalent bonds instead of the expected two?
Carbon undergoes electron promotion, where one electron from the 2s orbital is promoted to the empty 2p orbital, giving carbon four unpaired electrons. This allows carbon to form four covalent bonds (e.g.,
C
H
4
CH
4
).
Why is electron promotion in carbon energetically favorable?
The energy required for electron promotion is small and is more than compensated by the energy released when forming four bonds, making it favorable for carbon to form four bonds instead of two.
How does phosphorus expand its octet to form five covalent bonds?
Phosphorus can promote an electron from the 3s to the empty 3d orbital, allowing it to have five unpaired electrons. This enables phosphorus to form compounds like
P
C
l
5
PCl
5
with five bonds.
Why can phosphorus form
P
C
l
5
PCl
5
but not
P
I
5
PI
5
?
The
P
−
I
P−I bond is weaker than the
P
−
C
l
P−Cl bond because iodine is larger and forms weaker bonds with phosphorus, releasing less energy, which is insufficient to compensate for the energy required to promote an electron.
Describe the difference between a single bond, double bond, and triple bond in terms of sigma (σ) and pi (π) bonds.
Single bond: Consists of one sigma (σ) bond from head-on orbital overlap.
Double bond: Consists of one sigma (σ) and one pi (π) bond; the pi bond is formed by sideways overlap of p orbitals.
Triple bond: Consists of one sigma (σ) and two pi (π) bonds.
Explain the concept of a sigma (σ) bond and a pi (π) bond.
A sigma (σ) bond forms through head-on overlap of orbitals directly between two nuclei, while a pi (π) bond forms from the sideways overlap of p orbitals above and below the bond axis, only occurring in double or triple bonds.
Why is a pi (π) bond generally weaker than a sigma (σ) bond?
A pi (π) bond is weaker because it involves the sideways overlap of p orbitals, resulting in electron density that is less directly between the nuclei compared to the head-on overlap in sigma (σ) bonds.
Can atoms rotate freely around a double or triple bond? Why or why not?
No, atoms cannot rotate freely around double or triple bonds because the pi (π) bonds restrict rotation. Breaking a pi bond would be required to allow rotation, which requires energy.
No, atoms cannot rotate freely around double or triple bonds because the pi (π) bonds restrict rotation. Breaking a pi bond would be required to allow rotation, which requires energy.
At high temperatures, the pi (π) bond may temporarily break, allowing rotation around the double bond. Once cooled, the structure typically returns to its original state with the pi bond intact.
Which bonds allow rotation
Single as only made of a sigma bond,but not double or triple as pi bond restricts rotation
Why are geometric isomers stable at room temperature
Geometric isomers have restricted rotation around double bonds due to pi (π) bonding, making the arrangement of atoms fixed and stable at room temperature.
How is covalent bond strength measured?
Covalent bond strength is measured by the amount of energy required to break the bond in a gaseous molecule.
What factors influence covalent bond strength?
The sum of the atomic radii of the bonded atoms, Smaller atoms and multiple shared electron pairs lead to stronger bonds.
Why does diamond have a high melting point?
Diamond has a giant covalent structure where each carbon atom is covalently bonded to four others in a 3D lattice, requiring a large amount of energy to break all covalent bonds.
Describe the bonding in silicon dioxide (SiO₂).
Each silicon atom is covalently bonded to four oxygen atoms, and each oxygen atom is covalently bonded to two silicon atoms, forming a 3D lattice similar to diamond.
How do the melting points of diamond and SiO₂ compare, and why?
Diamond has a higher melting point than SiO₂ due to the strong C-C bonds in its structure, which are stronger than the Si-O bonds in silicon dioxide.
Why is the F-F bond weaker than expected for such a small atom?
The F-F bond is weaker than expected because the small size of fluorine atoms brings lone pairs close together, leading to significant repulsion between these electrons. This repulsion weakens the bond, making F-F less strong compared to other single bonds between small atoms.
What is bond energy?
Bond energy is the amount of energy required to break one mole of a bond in a gaseous molecule. It reflects the bond’s strength and is influenced by factors like the size of the atoms and the number of shared electron pairs.
What determines the strength of ion-ion interactions in ionic compounds?
The strength depends on the charge of the ions (higher charge = stronger attraction) and the ionic radius (smaller ions = stronger attraction). These are non-directional forces that contribute to the reason ionic compounds form crystal lattice structures as there and no specific angles
How does bond length affect the strength of covalent bonds?
Shorter bonds are generally stronger because the nuclei of the atoms are closer, increasing electrostatic attraction. Longer bonds are weaker due to increased distance.
Explain why smaller metal ions form stronger bonds in a metallic structure compared to larger metal ions.
Smaller ions have more delocalised electrons which is closer to positive charged nucleus ,more compact lattice so electrostatic force of attraction is greater .
Describe ion-dipole interactions and provide an example where this interaction is significant.
Ion-dipole interactions occur between a charged ion and a polar molecule. For example, the attraction between Na
+
+
ions and water molecules is crucial for dissolving salts in water.
What causes permanent dipole-dipole forces, and between which types of molecules do they occur?
They are intermolecular forces that occur between polar molecules as they contain atoms of different electronegativities which leads to a distribution of electrons densities
How do London forces arise, and why are they stronger in larger atoms or molecules?
London forces arise from temporary charge imbalances within a molecule, creating an instantaneous dipole that induces a dipole in a nearby molecule. They are stronger in larger atoms or molecules because they have more electrons, which are more easily polarized.
Explain ion-water (polar) interactions
Water is a polar molecule positive metal ions would be attracted to slightly negative oxygen in water as one of the lone pairs in oxygen forms a dative covalent bond with empty orbital in metal. and negative ions would be attracted to positive hydrogen in water ,the water makes ions hydrated
What are London forces?
London forces are weak, temporary attractions between molecules, caused by constantly shifting temporary dipoles due to random electron movement.
How does the number of electrons in a molecule affect the strength of London forces?
Molecules with more electrons have stronger London forces, as seen in iodine (solid) compared to chlorine (gas) at room temperature.
Why does it take more energy to overcome dipole-dipole forces compared to London forces?
Dipole-dipole interactions are generally stronger because they are permanent, while London forces are temporary and weaker.
How does molecular shape and contact points influence London forces
Molecules with more contact points experience stronger London forces. For example, long molecules like pentane have stronger forces than spherical molecules like neopentane.
How does molecular shape and contact points influence London forces
Molecules with more contact points experience stronger London forces. For example, long molecules like pentane have stronger forces than spherical molecules like neopentane.
Why do noble gases have low boiling and melting points?
Noble gases have fewer electrons and lack bonding electrons, resulting in very weak London forces between atoms.
What is hydrogen bonding and what elements are typically involved?
Hydrogen bonding is a strong type of intermolecular force, occurring between a δ⁺ hydrogen in one molecule and a lone pair of electrons on a small electronegative atom (usually fluorine, oxygen, or nitrogen) in another molecule.
Why can’t hydrogen bond with chloride ions despite both having δ⁻ regions?
Chloride ions are too large, which makes hydrogen bonds ineffective, even though chlorine has a δ⁻ charge.
Why does water (H₂O) have a higher boiling point than hydrogen fluoride (HF)?
Water forms more hydrogen bonds per molecule due to two lone pairs on oxygen, making it stronger overall in terms of hydrogen bonding than HF
What happens when we heat a substance that has covalent bonds
Molecules move faster and intermolecular forces may break,its not the covalent bond that breaks, as intermolecular forces are much weaker than covalent bonds
When will an ionic compound conduct electricity
When molten/liquid or when dissolved in water
When talking about metallic bonds or ionic bonds what do you have to make sure to mention in questions
Ionic radius , charge on cation
What is electronegativity
The power of an atom to attract the shared pair of electrons in a covalent bond
What is the most electronegative element
Fluorine (4.0), small radius
Define bond polarity
When 2 atoms in a covalent bond have different electronegativities the bond electrons are attracted to the atom with a higher electronegativity creating a polar bond , this results in the more electronegative atom becoming delta negative and the less electronegative becomes delta positive
What does a molecules polarity depend on
It’s 3D structure and the polarity of its bonds
Define dipole moment
A measure of positive and negative charges in a molecule
Why do some molecules that have polar bonds are known as non polar overall
Their dipole moments cancel out so the molecule overall is non polar
How can i test for polarity
Run liquid near a charged rod or balloon ,polar liquids are attracted to charge as they have a dipole moment,non polar liquids are not effected
I placed an unknown substance by a charged rod and the substance was attracted to the rod ,what can u tell me about the substance
It is polar as it reacts to charge due to its dipole moment
What must the difference in electronegativity be for non polar covalent bonds,polar covalent bonds and ionic bonds
Non polar covalent, 0-0.4
Polar covalent, 0.5-1.7
Ionic, 1.7 and above
Why won’t a covalent bond with a difference in electronegativity of 0.4 dissolve in water?
Because its non polar covalent so won’t dissolve in water which is polar
The difference in electronegativity for a bond is 1.8 ,will this molecule dissolve in water?explain answer.
Yes as this is an ionic bond, and ionic bonds dissolve in water as water is polar and interacts with the positive and negative charges present in ionic bonds causing ionic bonds to dissolve in water
Define polymeric
Long chains of molecules bonded covalently with weak forces between chains
