Chemistry Fundamentals

Atomic Models & The Periodic Table

• Understanding the periodic table's organization simplifies understanding other concepts.
• Electrons are negatively charged and surround the positively charged nucleus.

Models in Science

• A model is a descriptive and predictive representation of a phenomenon.
• Effective models both explain and predict.
• Models are built from data and existing theories.
• Models evolve as new insights are gained and false hypotheses are disproven.
• Multiple models can be correct for a single phenomenon, each highlighting different aspects.
• Models apply to both microscopic and macroscopic scales, with insights transferable between them.
• Science is primarily a process of discovery and refinement, not just the final outcome.

Development of Atomic Theory

• Early atomic theory proposed indivisible particles as the basic unit of all matter.
• Atoms of one element cannot be converted into atoms of another element.
• Atoms of a given element are identical in mass and properties, differing from atoms of other elements.
• Compounds are formed through the chemical combination of atoms in specific ratios.
• Atoms consist of a nucleus containing protons and neutrons, surrounded by electrons.
• Protons have a positive charge, neutrons are neutral, and electrons are negatively charged.
• The number of protons and electrons determines an element's identity.
• Protons and neutrons have significantly higher mass than electrons.
• The mass of an atom is primarily determined by the number of protons and neutrons.

Counting Atoms by Mass (Moles)

• Carbon-12 serves as the standard for measuring atomic mass.
• A mole is equal to 6.02 x 10^23 units (Avogadro's number).
• One mole corresponds to the number of carbon atoms in 12g of Carbon-12.
• A mole can represent an element, compound, or mixture.
• An element consists of only one type of atom.
• A compound is a molecule composed of two or more elements.
• A mixture contains elements and compounds.

The Periodic Table: Organization

• Columns are called groups (18 in total).
• Rows are called periods (7 in total).
• Elements with similar properties are grouped by color.
• Elements are organized by type (metals, metalloids, nonmetals).
• Group 1: Alkali metals.
• Group 2: Alkaline earth metals.
• Groups 3-12: Transition metals.
• Group 18: Noble gases (inert).
• Group 17: Halogens.
• Differences between atoms/elements are based on the number of protons and electrons.
• Carbon-12: 6 protons, 6 neutrons, 6 electrons.
• Neutrons do not determine element identity; protons and electrons do.
• Element card patterns include appearance description (top right).
• Atomic mass unit (middle number).
• Reactivity (bottom left).
• Usual state (bottom right).
• Spokes indicate conservation with other elements that have the same spokes.

Exponent Rules

• To divide exponents with the same base, subtract the exponents: A^7 / A^4 = A^3.
• Negative exponent: A^-x = 1 / A^x.
• To multiply exponents with the same base, add them: V^3 * V^6 = V^9.

Bohr's Atomic Model

• Protons and neutrons are in the nucleus, with electrons surrounding it.

General Features of Atoms

• Atoms are electrically neutral, spherical, composed of a positively charged nucleus surrounded by negatively charged electrons.

Periodic Table Features

• Atomic number: Number of protons and electrons.
• Average atomic mass: Average mass of all isotopes, weighted by abundance.
• Isotopes with a presented mass that is closer to a known element's atomic mass are more abundant.
• To find the number of neutrons in an isotope, subtract the atomic number from the isotope's mass number: e.g., 192Pk (78 protons) has 192 - 78 = 114 neutrons.

Calculating Average Atomic Mass

• Sum the mass of each isotope.

Periodic Table: Predictions

• Can predict if a compound is ionic (+/-) or molecular.
• Predicts ions formed by main group elements (not transition metals).
• Predicts formula/name of ionic compounds.
• Predicts some polyatomic ions.

Periodic Table Insights

• Provides the number of protons and electrons.
• Metals lose electrons to become cations (positive ions).
• Group 1 elements lose one electron to become +1 cations.
• Group 2 elements lose two electrons to become +2 cations.

Compound Classes

• Ionic Compounds
  • Formed between a metal and a nonmetal.
  • Example: Table Salt (NaCl).
  • Tend to be solids at room temperature.
  • Have high melting points.
  • Brittle.
  • Good conductors of electricity.
• Covalent/Molecular Compounds
  • Formed between two nonmetals, or a nonmetal and a metalloid.
  • Example: Water (H2O).
  • Usually gas or liquid at room temperature.
  • Have low melting points.
  • Malleable
  • Poor conductors.
• Left side of periodic table is mainly metals (except hydrogen).
• Right side is mainly nonmetals.
• Group 18 elements (noble gases) are nonreactive.
• Other nonmetals tend to gain electrons to achieve noble gas electron configurations.
• Group 17 elements gain one electron to become anions.
• Periodic table is arranged by atomic number.
• Atomic number indicates the number of electrons and protons.
• Group 1 elements have 1 valence electron (available for bonding).
• Atoms aim to have 8 valence electrons in their outermost shell (octet rule).
• Metals lose electrons to become cations.

Metals vs. Nonmetals

• Metals conduct electricity, are malleable, lose electrons (become cations), and react with nonmetals.

Naming Ions

• Metal name followed by nonmetal name with "-ide" suffix
• NaCl: Sodium Chloride
• KBr: Potassium Bromide

Isotopes

• Atoms with the same number of protons but different numbers of neutrons.

Mass Defect

• When protons and neutrons combine in the nucleus, energy is released, leading to a mass defect.
• The mass of an atom is always less than the sum of its component particles.

Bonding Patterns

• Ionic compound: metal and nonmetal, with ionic bonding.
• Molecular compound: two nonmetals or a nonmetal and metalloid, with covalent bonding.

Monoatomic Ion Charge

Determined by their position on the periodic table.

• Group 2 elements become +2 cations (lose electrons).
• Nonmetals in Groups 15-17 become negative anions (gain electrons).

Shell Model of the Atom

• Classifies electrons as core or valence.
• Lithium (2nd period, 1st group) has 2 shells and 1 valence electron.
• First shell holds two electrons maximum.
• Second shell holds eight electrons maximum.
• Number of electron shells = period number.
• Number of valence electrons = group number (for groups 1-2 and 13-18).

Determining Core Electrons

• Example: Sulfur (atomic number 16, group 16) has 6 valence electrons, thus 10 core electrons.
• Valence electrons are arranged predictably; core electrons are more complex.

Shell Models Across a Period

• The number of valence electrons increases across a period (row).

Interparticle Forces

• Closer particles = stronger attraction/repulsion (potential energy).
• Attractive forces are negative.
• Repulsive forces are positive.
• Factors: charge state, distance, magnitude of charges.

Coulomb's Law

• Force proportional to charge magnitude.
• Force inversely proportional to the distance between charges.
• Attraction increases across a period as charge increases.
• Attraction increases up a group as size (number of shells) decreases.
• Electrons are held in place by Coulomb's Law.

Periodic Table Trends

• Effective Nuclear Charge (Zeff)
  • The pull of the nucleus experienced by valence electrons.
  • Attraction decreases as distance increases.
  • Zeff = total electrons - core electrons.
  • Zeff increases across a period (row).
  • Zeff remains relatively constant down a group but may decrease slightly with added shells.
• Mass
  • Increases across a period due to added protons and neutrons.
  • Increases substantially down a group due to added shells.
• Atomic Radius
  • Increases down a group as shells are added to reduce electron-electron repulsion.
  • Decreases across a period due to increased attraction between electrons and the nucleus.
• Zeff and Atomic Radius have an inverse relationship.
• Cations are smaller than their corresponding neutral atoms (less electron-electron repulsion).
• Anions are larger than their corresponding neutral atoms (more electron-electron repulsion).
• Atomic size decreases as Zeff increases.

Reactivity

• Metals are more reactive.
• As you go down the group elements become more reactive
• Nonmetals are more reactive on the right side of the periodic table.
• Going down the group nonmetals are less reactive
• Elements at the top of the table have higher Zeff so they want to gain an electron

Polarity & The Octet Rule

• All atoms seek to achieve a stable octet (8 valence electrons).

Ionization Energy

• The energy required to remove the outermost electron from an atom.
• Higher energy is required to remote a core electron.

Electronegativity

• The tendency of an atom to attract electrons in a bond.
• Electronegativity values indicate bond polarity.
• Values generally increase from bottom left to top right on the periodic table.
• They are always positive.

Polarity

• Separation of charge
• A molecule possesses polarity when one side is distinctly different from the other.
• Carbon-hydrogen bond is the threshold between polar and nonpolar.
• Oils are nonpolar due to primarily carbon-hydrogen bonds (hydrocarbons).
• Electronegativity and polarity are directly related.
• Electronegativity difference of 0 > 0.4 = Nonpolar.
• Electronegativity difference of 0.5 > 1.7 = Polar.

Polar/Nonpolar Bonds

• Common polar bonds: Carbon-oxygen, nitrogen-hydrogen, hydrogen-fluorine, bonds with bromine or chlorine.
• Common nonpolar bonds: Carbon-carbon, carbon-hydrogen, bonds between two identical atoms.
• Electronegativity in molecules creates dipole moments (regions of partial charge).

Depicting Polar Bonds

• Unequal sharing of electrons is shown with a polar arrow (arrowhead points to the more electronegative element).
• Partial charges can be indicated with delta plus (δ+) and delta minus (δ-) symbols.
• Delta negative will be on the more electronegative side.

Ionic Models

• Complete transfer of electrons creates positive (cations) and negative (anions) ions when electronegativity difference is > 1.7.
• Covalent bonds share electrons (unevenly in polar covalent bonds).
• Ionic bonds completely transfer electrons.
• Pure Ionic
  • No sharing of electrons
• Polar covalent
  • Bonding electron pairs shared unequally
• Pure covalent bonds
  • Bonding electron pairs shared equally

Net Dipole Moments

• Dipole moments can exist across axes of a molecule or can cancel out, resulting in a net dipole moment of zero.

Bonding & Bond Patterns

• Covalent and Ionic bonds
  • Similarities and differences between ionic and covalent bonds and what determines the type of bond form
• Octet Rule: Eight valence electrons is stable = Noble Gas Configuration.
• Atoms tend to gain or lose electrons to achieve this configuration.

Electron Transfer

• Example: MgCl2
• Mg bonds with two Cl because Mg has a +2 charge and Cl has a -1 charge which effectively fills both octets.

Ionic Solid Hardness

• The charge states are constant between Mg and Ca but since Ca has a greater atomic radius it cannot bond as tightly with oxygen, thus MgO is a harder ionic compound

Covalent Bonds

• Formed between two nonmetals or a nonmetal and a metalloid.
• Left over valence electrons are lone pairs (located on the most electronegative atom).

Bond Order

• Single bond = bond order 1.
• Double bond = bond order 2.
• Triple bond = bond order 3.
• Count valence electrons for accurate model building.

Bond Lengths

• Bond length decreases as the bond order increases (single to double to triple bonds).
• Bond Strength
• The shorter the covalent bond the stronger it is (when changing one variable)
• The more shared electrons shared between atoms, the stronger the bond (higher bond order = stronger bond)

Bonding Energetics

• Bonding lowers the energy of the system, making it favorable and more stable.

Lewis Structures

• Models that use dots and lines to represent valence electrons and bonds.

Formal Charge & Lewis Structures Cont.

• Atoms in molecules tend to share 2, 4, or 6 electrons to be stable = lower energy state.
• There are trends in bond # formation with PT group #
• Elements in group 7 only like to form one pond
• Elements in group 6 need two electrons so they like to make two bonds

Common Molecular Groups

• Carbon double bond with Oxygen (has two lone pairs)
• Carbon oxygen (two lone pairs) and carbon single bonded
• Carbon Oxygen (two long pairs) Hydrogen single bonded
• Carbon Nitrogen (a lone pair) and two hydrogen atoms single-bonded
• Carbon double bonds to nitrogen (a lone pair) and one hydrogen atom

HONC Rule

• Hydrogens like to form one bond
• Oxygens like to form two bonds
• Nitrogens like to form three bonds
• Carbon like to form four bonds
• Drawing C2H5NO

Building upon Lewis structure using formal charge

- We’re always trying to reach the lowest formal charge state on an atom and for a given molecule as a whole

• Formal charge is
  • The valence electrons on the atom of interest minus half of the electron assigned to others, minus all of the lone pairs assigned to the atom of interest

Follow these rules

 - Assign all electrons of a lone pair to the atoms it’s on
 - Assign half the electrons of a bonding pair to the atom it’s one

• Choosing the “better” structure
  • Smaller formal charges (either positive or negative) are preferable to larger charges
  • Try to avoid like charges on adjacent atoms
  • A more negative formal charge should exist on a more electronegative atom

Lewis dot structures when there is an ion

 - With ions the HONC rule is not always the best option
 - If you have an anion you should add an electron
      - Add the extra electron to the most electronegative element
 - If you have a cation you have to remove an electron
      - With anions and cations you might not fill an octet
 - Put the charge associated with the compound on a bracket around the compound
      -  Ex: H30+

Hyper valence

• Compounds with a central atom in period 3 or higher can pair their lone electron pairs with others, this will break the octet rule

Productively breaking the rules

• The Lewis Structure Expansion Pack
• Beyond HONC 1234
• Valence shell expansions: Atoms of elements in period 3 and higher can sometimes fit more than 4 pairs of valence electrons around them
  • Ex: PCl5
  • Ex; Br3-
• The right answer
  • Just toss electrons wherever they fit best

Situation 2

• Electron deficiency: There are not enough valence electrons to fill all octets; these are often reactive molecules, often cations
  • Ex: Ch3+
• You know there's an electron deficiency if there is no way to fulfill an octet

Situation 3

• Electropositive element exception: Some atoms, namely B, Al, and Be, can form compounds with less than 4 bonds
  • Ex: BF3

Situation 4

• Odd electron exception: a few molecules have an odd number of electrons that means it is necessary to put a single electron somewhere
  • Ex: NO

Thought process when building elements

 - Which molecule is more electropositive? Place in the middle
 - How many valence electrons do I need to account for
 - How many bonds does my central atom need to fill the octet
 - Can I expand the valence shell
 - Do I have an electron-deficient molecule? Cations
 - Electropositive elements? B, Be, A
 - Odd # of electrons?

Resonance

Occurs when more than one valid Lewis structure can be written. The actual structure is an average or superposition of the three structures

• Delocalized electrons will be drawn as dashed lines
• Backbone of resonance structures will always be the same whereas isomers can have different backbone structures (only the arrangement of electrons is different)
• Drawing C3H7Br

Resonance & Intro to Molecular Geometry

• Resonance occurs when more than one valid Lewis structure can be written. The actual structure is an average / superposition of all three structures
• Resonance is a system oscillating back and forth
• Delocalized oscillations of electrons - These electrons are moving back and forth from one place in a molecule to another without impacting the backbone or net formal charge Ex: Acetate CH3COO-
• Isomer: each of two or more compounds with the same formula but a different arrangement of atoms in the molecule (and different properties)
• Resonance structures only the electrons move
• A resonance structure is kind of like a hybrid molecule (arrows mean that each structure is equally as likely to occur)
  • EX: S3
• Rules you can productively break
  • Ions
  • Resonance structures
  • Valence shell expansion
  • Electron deficiency
  • Electropositive element
    • Odd electron exception

Introduction to Molecular Geometry

-We can predict molecular geometry with VSPER: valence shell electron pair repulsion -If we model the molecule of interest the shapes will place the electrons as far away as possible from each other -Lone pair electrons are placed in front and behind the nucleus -Any given central atom can have 360 degrees worth of possible bonding -Only bonded atoms name the geometry Meaning lone pairs are not accounted for when naming the geometry

Linear

Ex: CO2 has a 180-degree bond angle

Trigonal Planar

Ex: BF3 has a 120-degree bond angle -Structures lessen e- repulsion to the ideal angle

Trigonal Pyramidal

Ex: NH3 has a lone pair on top; that is why the “legs” are pointing down Approximately 107 degrees apart

Tetrahedral arrangement

• Ex: CH4 -Straight lines are in view -Filled in triangle is towards me -Perpendicular lines are going away from me
• Has a ideal angle of around 109.5 degrees

Trigonal bipyramidal

Ex: PCL5

-Middle axis is 120 degrees apart (trigonal planar) -Vertical axis is 180 degrees apart (Linear)

Octahedral

-Middle axis is 90 degrees apart

• Linear aspect in the vertical plane

Steps to take when determining molecular geometry

-NH3 - Draw lewis structure - Count the electron pairs and put them as far away as possible - Determine the position of the atoms based on electron pairs and occupancy - Name the structural form’ Trying to draw CS2

Electron Pair Geometry

-The arrangement of the electrons both bonded and lone pairs around a central atom -For this we treat the atoms and electrons equally -You can understand what shape you will get if you calculate the steric number of a compound

of lone pairs + # of total atoms = steric number

 - 2 = linear
 - 3 = trigonal
 - 4 = tetrahedral

Molecular Geometry Continued

• Bent structures vary in bond angles because of their different back bones (depends on what atoms are bonded together)
• Bends and Pyramidal’s with a lone pair is less than ideal

Examples

• IF5 is square pyramidal
• SO2 is bent (root trigonal planar)
• SO3 2- is (base tetrahedral) Trigonal Pyramidal

Deviations from VSPER angles

• Bond pair - lone pair repulsion with make the bond angle of a tetrahedral molecule smaller
• Lp - lp repulsion makes the bond angle even smaller
• Higher-order bonding also makes angles smaller as they take up more space
• Lone pairs have a greater impact on angles when both lone pairs and higher-order bonds are present
• Double and triple bonds with resonance Since the electrons are delocalized in a molecule with a resonance structure, then the bond angle is still ideal because the electrons can go everywhere
• Geometry and Polarity
• NH3 vs CF4
• In terms of resonance structures since the electrons are delocalized there is no place the electrons reside so they are non polar

Intermolecular Farces (IMF)

• When molecules have polarity, they can act in ways that are repulsive or attractive
• What defines a nonpolar molecule
  • Something that is composed of bonds where the difference in electronegativity is 0.4 or less
• Nonpolar molecules can have lightly attractive or repulsive forces in brief and fleeting ways
• Understanding kinetic energy ->Kinetic energy is once something starts interacting or moving about > Energy primarily comes from heat -The more heat you put into a system the more energy you put into it
• Once we consider the total energy of particles, we have to consider the mass of the molecules as well -> Moving bigger particles is harder (takes more force) -> E = ½ mv^2
• When molecules interact with each other we call this intermolecular force
  • This is in opposition to an intramolecular force (a force within something)
• Covalent forces - Unequal sharing of electrons
• Ionic forces - Giving and receiving electrons
• Metallic bonds
  • Metals like to lose electrons to become cations
    • All atoms in a sodium body give out their electrons and they have a positive charge around them
• Intermolecular forces arise when mixtures or solutions of molecules collide and interact due to kinetic energy

Typers of intermolecular forces

• London - dispersion forces or Induced dipole - The weakest IMF (intermolecular force) Electrons move around their space and it is favorable to move all electrons to one side ->This creates a slightly negative and positive side -> Creates a very fleeting dipole moment due to energetic favorability -London dispersion is defined by Van Der Waals radii - A balance of creating a favorable IMF vs the penalty of placing nuclei at a repulsive distance
• More polarizable molecules create strong london dispersion - Forces - more electrons = more ability to polarize Lower in period means more electrons which means more ability to polarize Ex: Can be anything
• Dipole-dipole interactions -Two or more molecules that have dipole moments can interact with each other - The negative side will bond with the positive side of another molecule -> Stronger than LDF VS LDF ->Dp dp are based on permanent dipole opposed to LDF’s induced dipoles ->Dp dp are stronger in force than LDF which are weaker

Ex: CH3Cl

• H-bond interactions -Happens when a hydrogen atom is in close proximity to a nitrogen, fluorine, or oxygen atom -Because of hydrogen’s electron deficiency it is stronger than dipole interactions’ H-bond donors and H-bond accecptors -Donors are attached directly to the hydrogen atom -EX: H20 the donor is the oxygen - The acceptor is the electronegative atom in the bond -H-bonds can exist within mixtures So strong that they have somewhat covalent characteristics, hard to break apart EX: H20
• Ion - dipole interactions -The strongest of them all - One partner in the IMF is a permanent ionic state the other is a permanent dipole -The magnitude of the charges and the dipole have a direct relationship with the strength of bonds -Permanent ions tend to be attracted to a permanent dipole -EX: Salt water
• Polar: H-bonds, dipole dipole, ion-dipole -Non polar: London dispersion

IMF & Solubility

• Solutions can be called hydrophillic of hydrophobic-Hydrophilic: H-bond, dipole dipoole, ion-dipole Are all water soluble -Hydrophilic: London dispersion forces
• Intermolecular forces play a role in dinner -Like dissolved like -If two molecules are soluble, they mix with minimal penalty -Olive oil and vinegar don’t mix because the oil is not soluble in vinegar -Do not have similar properties -Do not have same IMF -Energetic penalty to mix
• Vinegar and water are soluble because water can participate in h-bonds, dipole dipole and ion dipole - Vinegar has h bond acceptor and h bond donor -Since they both have these properties they are soluble in each other
• Polar molecules
• Nonpolar molecules -Bond dipoles cancel -No net dipole moment -Soluble in nonpolar solvents
• Understanding kinetic energy -Speed, mass, and kinetic energy are also directly related -Ek = ½ mv^2 -As the temperature increases so does the kinetic energy -Ek = 3/2 KbT >KbT = Boltzman constant >Puts value to the energy associated with the random motion of molecules -Potential energy is the energy of attraction -Coulombs law: depends on the radius and charge of the atom or whatever
• Molecules at the same temperature have the same average kinetic energy-Heavy molecules move more slowly-Lighter molecules move faster

States of matter

 -Solids and Liquids and Gasses -Solid: tightly packed molecules
      -If we put energy into the system

• It becomes liquid and eventually a gas -Liquid isn’t as tightly packed as ice -Gas is even more erratic than liquid
• Solubility Cont. & Scientific Measurement
• Intermolecular forces can be split up into polar or nonpolar-Polar: Hydrogen bonds, dipole-dipole, and ion-dipole-Nonpolar: London dispersion forces -All molecules can participate in ldf but not everything can participate in polar bonds
• Hydrophilic or hydrophobic ->Hydrophilic molecules are largely polar ->Hydrophobic molecules are largely nonpolar
• Salad dressing is made up of Nonpolar oils >Hydrocarbon rich molecules Polar layer is vinegar and water

Solids Liquids and Gases

• Solid Fixed volume, does not take shape of container -Water is solid at 263K -Liquid -Fixed volume, takes shape of container -Water is liquid at 291K -Gas
• Kinetic molecular view of the three states -Gas; Attractive forces are weak relative to kinetic energy >Liquid: Attractive forces are stronger because particles have less kinetic energy >Solid: -The boiling point is the highest temperature at which liquid exists for a specific atmospheric pressure
• Boiling point vs. shape-The larger area of contact between atoms means a higher boiling point -Boiling points vs number of electrons

As you add shells of electrons you add electrons that can participate in -LDF which means it will be stronger (boiling point will be higher)

• Boiling points for compounds with hydrogen >The energy of the h-bond (H2O) is stronger than LDF therefore takes more energy to break apart (to boil) The more electrons something has the harder it is to boil
• What impacts boiling point >Polar vs. Nonpolar: Higher polarity means higher boiling point>Chain Length: Longer the chain, higher boiling point>Surface area of molecule (number of branches): Higher area, higher boiling point>Molar Mass: Higher molar mass, higher boiling point

Mixture vs. Solution

• Mixture: insoluble-Ex: oil and vinegar-Solution: Soluble-Ex: blue dye in H20 -Definitions of words relating to solutions-Solvent: the liquid-Solute: The substances that dissolves in the liquid (could be a gas, liquid, or solid)-Solution: The homogeneous mixture of a solute and solvent -Why do polar substances tend to dissolve in water -Dipole-dipole-attractions and hydrogen bonding When representing something dissolved in water you write (aq) after the molecular formula, this means aqueous, which is dissolved in water Like dissolves like If attractions between two types of molecules are similar in strength they will mix>Energetically favorable If attractions between two types of molecules are different in strength they will not mix>Energetic penalty Solubility of Alcohols At what point do LDF outweigh the H-bonds to influence the solubility of alcohol At around 5 carbons the solubility drastically changes Cell membranes The head is polar so its hydrophilic The tails are hydrophobic because they are made of carbon hydrogen chains What is scientific notation

A way of representing numbers that are really big or really small Glucose

• Has a molar mass of 180.56g per mole -Molar mass is the sum of the molecules that make up a molecule Scientific notation >You have a coefficient of whatever, multiplied by base 10 and an exponent>The coefficient has to be between 0 and 9.999…. > The base will always be 10 -The exponent will be some positive or negative integer Exponents with positive values are greater than 0 Exponents with negative values are less than 0p Significant figures Any number that is nonzero unless the zeros are in between two nonzero integers S.I. Units & Math (last info for 2nd midterm, will need a calculator and pencil, very math dense)
• Dimensional Analysis with prefixes Just cross out your units no matter how many this equals that there is -How to move numbers around >100g of sugar could be thought of as a single hectogram of sugar How does the rest of the world read temperature Celsius, 0 is when water freezes, Kelvin 0 is when nothing moves not even electrons What is a Dalton? -1 dalton is equal to one gram per mole ( 1g/mol) or atomic mass unit (amu)

A mole

Is a unit of measurement is based on the mass of 12 grams of carbon Also 6.02 x 10^23 atoms of something -Going from grams to mols If you have 2 grams of NaCl how many mols do I have Grams per mols also applies to liquid - grams is the S.I. bas unit for mass Molarity of a liquid = mols of solute / volume of solute >10M = 10 mols of NaCl / 1 L of H2O

Test your knowledge of fundamental chemistry concepts. Questions cover atomic theory, moles, elements, the periodic table, isotopes, and intermolecular forces. This quiz is designed to assess understanding of basic chemical principles.