S2 + S3 Flashcards
(173 cards)
1
Q
when metal atoms lose electrons, they form
A
positive ions called cations
2
Q
when non-metal atoms gain electrons, they form
A
negative ions called anions
3
Q
the ionic bond is formed by
A
electrostatic attractions between oppositely charged ions
4
Q
binary ionic compounds are named
A
with the cation first followed by the anion. The anion adopts the suffix ‘ide
5
Q
ionic compounds exist as
A
three-dimensional lattice structures, represented by empirical formulas
6
Q
lattice enthalpy is a measure of
A
the strength of the ionic bond in different compounds, influenced by ion radius and
charge.
7
Q
ionic compound melting point
A
Melting an ionic compound to form a liquid involves breaking the electrostatic attractive forces between the oppositely charged ions in the lattice. Due to the strength and number of the electrostatic attractive forces between ions, a lot of energy is required to melt ionic compounds. Therefore, ionic compounds have high melting points.
8
Q
ionic compounds volatility
A
As ionic compounds have high melting points and do not easily vaporise they have a low volatility.
9
Q
ionic compounds solubility
A
Many, but not all, ionic compounds are soluble in water. The solubility of an ionic compound depends on the strength of the attraction between the ions and water along with the strength of the ionic bond. An ionic compound is insoluble in water when it does not form a homogenous mixture when added to water.
10
Q
electrical conductivity ionic compounds
A
solid no
molten yes
11
Q
ionic compounds brittleness
A
oui
12
Q
a covalent bond is formed by
A
the electrostatic attraction between a shared pair of electrons and the positively charged nuclei
13
Q
the octet rule refers to
A
the tendency of atoms to gain a valence shell with a total of 8 electrons
14
Q
single, double and triple bonds involve
A
1, 2 or 3 shared pairs of electrons respectively
15
Q
a coordination bond is
A
a covalent bond in which both electrons of the shared pair originate from the same atom
16
Q
the valence shell electron-pair repulsion (VSEPR) model enables
A
the shapes of molecules to be predicted from the repulsion of electron domains around a central atom
17
Q
bond polarity results from
A
the difference in electronegativities of the bonded atoms
18
Q
molecular polarity depends on
A
both bond polarity and molecular geometry
19
Q
carbon and silicon form
A
covalent network structures
20
Q
the nature of the force that exists between molecules is determined by
A
size and polarity of the molecules
21
Q
intermolecular forces include
A
London (dispersion) forces, dipole-induced dipole, dipole-dipole and hydrogen bonding
22
Q
given comparable molar amss, the relative strengths of intermolecular forces are generally
A
London (dispersion) forces < dipole-dipole forces < hydrogen bonding
23
Q
chromatography is
A
a technique used to separate the components of a mixture based on their relative attractions involving intermolecular forces to mobile and stationary phases
24
Q
Bond strength
A
measure of the energy required to break a bond. As you may have guessed, the more bonds there are, the stronger the bond
25
electron domain geometry
26
polar nature of a covalent bond from electronegativity values
27
Diamond
Each carbon in diamond is covalently bonded to four other carbons.
28
properties of diamond
hard, high melting and boiling point, no conduct
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Graphite
Graphite contains carbon atoms which are joined to three other carbons.
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properties of graphite
soft and slipery, lubricant
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Graphene
composed of just one layer of graphite and is said to be the strongest, lightest and thinnest material known to humans. It is so thin, just one carbon atom thick, that it can technically be referred to as being two-dimensional.
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properties of graphene
strong, conduct (nanotubes),
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Fullerenes
similar to graphene but the carbon atoms connect together to form spherical shapes
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properties of fullerene
conduct (conductivity is much poorer than that of graphite. Although delocalised electrons are present, because they can only move around each fullerene and not jump between two separate fullerenes limits their conductivity because they cannot move in along a linear path as with graphite and graphene.),
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Silicon
Pure silicon contains one silicon atom that is covalently bonded to four other silicon atoms
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properties of silicon
electrical conductivity is poor, high melting and boiling point
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Silicon dioxide
Silicon dioxide, SiO2, is composed of a silicon bonded to four oxygen atoms with each oxygen also bonded to 2 silicon atoms. SiO2 is the simplest ratio
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properties of silicon dioxide
melting and boiling of silicon dioxide are very high, insulator
39
The types of intermolecular force are
London dispersion
Dipole-induced dipole
Dipole–dipole
Hydrogen bonding
40
London dispersion forces
the random, continuous movement of electrons can create asymmetrical electron distribution which causes attraction between neighbouring atoms or molecules.
41
Dipole–dipole
This type of force will only exist between two asymmetrical molecules that each have a permanent dipole. Dipole–dipole forces are much stronger than London dispersion forces, but can act in unison with them
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Dipole-induced-dipole
if a very polar molecule approached a molecule which is non-polar
43
hydrogen bonding
Hydrogen bonding is a type of intermolecular force that occurs when a hydrogen atom, covalently bonded to a highly electronegative atom (such as nitrogen, oxygen, or fluorine), interacts with a lone pair of electrons on another electronegative atom. This interaction is stronger than van der Waals forces but weaker than covalent or ionic bonds. Hydrogen bonds play a crucial role in the structure and properties of water, biological macromolecules like DNA and proteins, and many other substances.
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What is volatility in the context of covalent substances?
Tendency to evaporate easily.
46
What are the two types of covalent substances?
* Simple molecular
* Giant covalent
47
What characterizes simple molecular substances?
Weak intermolecular forces and low boiling and melting points.
48
Give two examples of simple molecular substances.
* Methane
* Water
49
What type of intermolecular forces are present in simple molecular substances?
* London dispersion forces
* Hydrogen bonding
50
What results in the high melting and boiling points of giant covalent structures?
Strong covalent bonds throughout the structure.
51
Provide two examples of giant covalent structures.
* Diamond
* Graphite
52
Do covalent substances generally conduct electricity?
No, because they do not have free-moving charged particles.
53
What is a notable exception to the electrical conductivity of covalent substances?
Graphite conducts electricity due to delocalized electrons.
54
Which molecular substances can conduct electricity in solution?
Molecular substances that ionize in solution, such as acids.
55
What affects the solubility of simple molecular substances?
Polarity.
56
Fill in the blank: Nonpolar molecules are soluble in _______ solvents.
nonpolar
57
Give an example of a nonpolar molecule.
Methane
58
Fill in the blank: Polar molecules are soluble in _______ solvents.
polar
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What is one reason giant covalent substances are generally insoluble?
Breaking strong covalent bonds requires too much energy.
60
What type of interaction allows polar molecules to be soluble in water?
Hydrogen bonding or dipole-dipole interactions.
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What is the volatility of simple molecular covalent compounds?
Volatile (low boiling/melting points)
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What is the volatility of giant covalent compounds?
Not volatile (high boiling/melting points)
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What is the solubility of polar covalent molecules?
Soluble in water.
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What is the solubility of nonpolar covalent molecules?
Insoluble/slightly soluble in water.
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What is the solubility of giant covalent structures?
Insoluble in water.
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What is the retardation factor (R_f) in chromatography?
A measure used in thin-layer chromatography (TLC) and paper chromatography to compare how far a substance moves relative to the solvent front.
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What is the formula for calculating R_f?
R_f = Distance traveled by the substance / Distance traveled by the solvent front.
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How is the distance traveled by the substance measured?
From the baseline to the center of the spot.
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How is the distance traveled by the solvent front measured?
From the baseline to the farthest point the solvent reaches.
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In the example given, what is the R_f value when a dye moves 4.5 cm and the solvent front moves 9.0 cm?
0.50.
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What does an R_f value closer to 0 indicate?
The substance is less soluble in the solvent and interacts more with the stationary phase.
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What does an R_f value closer to 1 indicate?
The substance is more soluble in the solvent and moves further up the plate.
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Why can different compounds be identified using R_f values?
Different compounds have different R_f values, allowing for identification by comparison with known substances.
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What factors can cause variation in R_f values?
Solvent polarity, stationary phase, and temperature.
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What are some uses of R_f values?
* Identifying unknown compounds by comparing to known standards
* Checking the purity of a substance
* Monitoring the progress of a chemical reaction.
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True or False: R_f values must be calculated under consistent conditions for accurate comparison.
True.
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Fill in the blank: R_f values range from ______ to ______.
0 to 1.
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metallic bond
electrostatic attraction between a lattice of cations and delocalised electrons
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the strength of the metallic bond depends on
the charge of the ions and the radius of the metal ion
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What property of metals allows them to conduct electricity efficiently?
Delocalized electrons that move freely throughout the metallic lattice
## Footnote
Metals are excellent electrical conductors because when a voltage is applied, these electrons flow easily.
83
Which metal is widely used in electrical wiring due to its high conductivity?
Copper (Cu)
## Footnote
Copper's high electrical conductivity makes it a preferred choice for wiring.
84
How do delocalized electrons contribute to thermal conductivity in metals?
They transfer thermal energy quickly by colliding and spreading the energy through the structure
## Footnote
This allows metals to conduct heat efficiently.
85
What is the role of tightly packed metal ions in thermal conductivity?
They conduct heat by passing vibrations (phonons) through the lattice
## Footnote
This enhances the overall thermal conductivity of metals.
86
Which metals are commonly used in cookware and heat exchangers due to their thermal conductivity?
Aluminum (Al) and Copper (Cu)
## Footnote
Their excellent thermal conductivity makes them ideal for these applications.
87
What defines the malleability of metals?
The ability of layers of metal atoms to slide over each other without breaking the metallic bond
## Footnote
This property allows metals to be hammered or rolled into sheets.
88
Which metals are known for their high malleability and can be shaped into thin sheets?
Gold (Au) and Silver (Ag)
## Footnote
These metals can be easily shaped, exemplified by gold leaf.
89
Fill in the blank: Metals exhibit unique physical properties due to their _______.
metallic bonding
## Footnote
This bonding involves positive metal ions surrounded by a sea of delocalized electrons.
90
What happens to delocalized electrons when a voltage is applied?
They flow easily, allowing the metal to conduct electricity efficiently
## Footnote
This is a key feature of metals' electrical conductivity.
91
True or False: Metals are poor conductors of heat.
False
## Footnote
Metals are known for their excellent thermal conductivity.
92
List two examples of metals that are excellent electrical conductors.
* Copper (Cu)
* Silver (Ag)
## Footnote
These metals allow electric current to flow efficiently.
93
List two examples of metals that are used for their thermal conductivity.
* Aluminum (Al)
* Copper (Cu)
## Footnote
These metals are preferred in applications requiring heat conduction.
94
What maintains cohesion in metals when the structure is deformed?
Delocalized electrons
## Footnote
They help maintain the integrity of the metallic bond during deformation.
95
What is the property of high electrical conductivity?
Free-moving delocalized electrons allow electric current to flow easily.
## Footnote
Uses: Electrical wiring, power cables, circuit boards. Examples: Copper (Cu), Silver (Ag).
96
What is the property of high thermal conductivity?
Heat spreads quickly through delocalized electrons and tightly packed atoms.
## Footnote
Uses: Cooking utensils, radiators, heat exchangers. Examples: Aluminum (Al), Copper (Cu).
97
What is the property of malleability?
Layers of atoms slide over each other without breaking bonds.
## Footnote
Uses: Car bodies, foil, metal sheets. Examples: Gold (Au), Aluminum (Al).
98
What is the property of ductility?
Metallic bonds allow metals to stretch without breaking.
## Footnote
Uses: Electrical cables, jewelry, metal mesh. Examples: Copper (Cu), Gold (Au).
99
What is the property of high density and strength?
Metals have closely packed atoms, making them strong and durable.
## Footnote
Uses: Construction materials, machinery, bridges. Examples: Iron (Fe), Titanium (Ti), Steel.
100
What is the property of shiny and lustrous appearance?
Metals reflect light due to free electrons.
## Footnote
Uses: Jewelry, mirrors, decorative items. Examples: Silver (Ag), Gold (Au), Aluminum (Al).
101
What is the property of corrosion resistance?
Some metals form protective oxide layers to prevent rusting.
## Footnote
Uses: Medical implants, outdoor structures, aircraft. Examples: Stainless Steel, Aluminum (Al), Titanium (Ti).
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What is the trend for melting points of s-Block Metals (Groups 1 & 2)?
Melting points decrease down the group as atomic size increases and metallic bonds weaken.
104
Why do Group 1 (Alkali Metals) have low melting points?
Weak metallic bonds due to one valence electron and large atomic radius.
## Footnote
Trend: Melting point decreases down the group (Li > Na > K > Rb > Cs). Example: Lithium (Li) has the highest MP in Group 1, Cesium (Cs) has the lowest.
105
What is the trend for melting points in Group 2 (Alkaline Earth Metals)?
Melting points are higher than Group 1 and decrease down the group.
## Footnote
Trend: (Be > Mg > Ca > Sr > Ba). Example: Beryllium (Be) has a much higher MP than Magnesium (Mg).
106
What is the trend for melting points of p-Block Metals (Groups 13-16)?
Melting points vary due to differences in bonding (metallic vs. covalent).
107
What is the trend for melting points in Group 13 metals?
Melting point generally decreases down the group, but Gallium has an unusually low MP due to weak metallic bonding.
## Footnote
Example: Aluminum (Al) melts at 660°C, Gallium (Ga) melts at just 30°C!
108
What is the trend for melting points in Group 14 metals?
Melting point decreases down the group as metallic bonding weakens.
## Footnote
Example: Tin (Sn) melts at 232°C, Lead (Pb) melts at 327°C.
109
What is the trend for melting points in Group 15 & 16 metals?
Low melting points due to weak metallic bonding.
## Footnote
Example: Bismuth (Bi) melts at 271°C, Polonium (Po) at 254°C.
110
bonding is best described as
a continuum between the ionic, covalent and metallic models, and can be represented by a bonding triangle
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the position of a compound in the bonding triangle is determined by
the relative contributions of the three bonding types to the overall bond
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alloys
are mixtures of a metal and other metals or non-metals; they have enhanced properties
113
polymers
are large molecules, or macromolecules, made from repeating subunits called monomers
114
addition polymers form by
the breaking of a double bond in each monomer
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What determines a material's properties?
The type of bonding determines the material’s properties.
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What are the properties of Ionic Bonding?
High melting/boiling points, conducts electricity in liquid/aqueous state, soluble in water.
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What are the properties of Covalent Bonding?
Low melting/boiling points, poor conductivity, varied solubility.
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What are the properties of Metallic Bonding?
Good conductor of heat & electricity, malleable & ductile, high melting/boiling points.
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How is the bonding type determined from electronegativity data?
Electronegativity difference (ΔEN) determines bonding type.
121
What are the steps to determine the position of a compound in the bonding triangle?
1. Find electronegativity values. 2. Calculate ΔEN. 3. Compare ΔEN to determine bonding type.
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What does ΔEN > 1.7 indicate?
Ionic Bonding.
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What does 0.4 < ΔEN < 1.7 indicate?
Polar Covalent Bonding.
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What does ΔEN < 0.4 indicate?
Non-Polar Covalent Bonding.
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What is an example of Ionic bonding?
NaCl (ΔEN = 2.1) → Ionic.
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What is an example of Polar Covalent bonding?
HCl (ΔEN = 0.9) → Polar Covalent.
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What is an example of Non-Polar Covalent bonding?
Cl2 (ΔEN = 0) → Non-Polar Covalent.
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What properties are predicted based on a compound's position in the bonding triangle?
Melting point, conductivity, and solubility.
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What are the properties of compounds with more Ionic character?
High melting/boiling points, soluble in water, conducts electricity in molten/aqueous state.
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What are the properties of compounds with more Covalent character?
Low melting/boiling points, does not conduct electricity, solubility depends on polarity.
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What are the properties of compounds with more Metallic character?
Conducts electricity and heat, malleable & ductile, insoluble in most solvents.
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Why are alloys stronger than pure metals?
Different-sized atoms disrupt the regular structure, preventing layers from sliding.
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What are the properties of alloys?
Less malleable, more durable, improved resistance to corrosion & wear.
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What is an example of an alloy and its properties?
Steel (Fe + C) → Stronger & rust-resistant.
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What is an example of brass and its properties?
Brass (Cu + Zn) → Harder than copper, used in instruments.
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What is an example of bronze and its properties?
Bronze (Cu + Sn) → Durable, used in statues.
137
What are the common properties of plastics?
Plastics are polymers with different bonding and molecular structures.
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What are the properties of Thermoplastics?
Weak intermolecular forces, soften when heated, can be reshaped & recycled.
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What are the properties of Thermosetting Plastics?
Strong cross-linked bonds, do not soften when heated, rigid & heat-resistant.
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What are the properties of Elastomers?
Flexible polymer chains, stretch & return to shape.
141
What is the process of drawing a repeating unit of an addition polymer?
Identify the monomer, break the double bond, extend the structure with brackets, add 'n' for repetition.
142
What is an example of polyethylene formation?
Monomer: CH2=CH2 (Ethene)
## Footnote
Polymer: −[CH2−CH2]−n (Polyethylene)
143
the periodic table consists of
periods, groups and blocks
144
the period number shows
the outer energy level that is occupied by electrons
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elements in a group have a common number
of valence electrons
146
periodicity refers to
trends in properties of elements across a period and down a group
147
trends in properties of elements down a group include
the increasing metallic character of group 1 elements and decreasing non-metallic character of group 17 elements
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metallic and non-metallic properties show a continuum; this includes the trend from
basic metal oxides through amphoteric to acidic non-metal oxides
149
the oxidation state is
a number assigned to an atom to show the number of electrons transferred in forming a bond; it is the charge that atom would have if the compound were composed of ions
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What is the structure of the periodic table?
The periodic table is organized into periods, groups, and blocks.
152
What are periods in the periodic table?
Periods are horizontal rows indicating the number of electron shells.
153
What are groups in the periodic table?
Groups are vertical columns where elements share the same number of valence electrons.
154
What are blocks in the periodic table?
Blocks classify elements based on their outermost electron subshell (s, p, d, f).
155
What does periodicity refer to?
Periodicity refers to trends in element properties across a period and down a group.
156
What happens to atomic radius across a period?
Atomic radius decreases from left to right due to stronger nuclear attraction.
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What happens to electronegativity across a period?
Electronegativity increases from left to right due to a stronger pull on shared electrons.
158
What happens to ionization energy across a period?
Ionization energy increases from left to right as more energy is needed to remove an electron.
159
What happens to metallic character across a period?
Metallic character decreases and non-metallic character increases from left to right.
160
What happens to atomic radius down a group?
Atomic radius increases from top to bottom due to more electron shells.
161
What happens to electronegativity down a group?
Electronegativity decreases from top to bottom due to a weaker pull on shared electrons.
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What happens to ionization energy down a group?
Ionization energy decreases from top to bottom, making it easier to remove electrons.
163
What happens to metallic character down a group?
Metallic character increases and non-metallic character decreases from top to bottom.
164
What are metallic properties?
Metallic properties include low ionization energy, good conductivity, and basic oxides.
165
What are non-metallic properties?
Non-metallic properties include high ionization energy, poor conductivity, and acidic oxides.
166
What are amphoteric elements?
Amphoteric elements show both metallic and non-metallic behavior.
167
What is the oxidation state?
The oxidation state represents the charge an atom would have if the compound were fully ionic.
168
What are the general oxidation state trends for metals?
Metals (Groups 1-3) have positive oxidation states as they lose electrons.
169
What are the general oxidation state trends for non-metals?
Non-metals (Groups 15-17) have negative oxidation states as they gain electrons.
170
What are the oxidation states of transition metals?
Transition metals have variable oxidation states (e.g., Fe²⁺, Fe³⁺).
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What are example oxidation states for NaCl?
In NaCl, Na has an oxidation state of +1 and Cl has -1.
172
What are example oxidation states for H₂O?
In H₂O, H has an oxidation state of +1 and O has -2.
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What are example oxidation states for Fe₂O₃?
In Fe₂O₃, Fe has an oxidation state of +3 and O has -2.
