Hoja de repaso: Fundamentals of Atomic and Molecular Chemistry

📋 Course Outline

1. Matter and states
2. Atomic structure
3. Atomic symbols and isotopes
4. Chemical formulae and naming
5. Periodic table trends
6. Bonding and electronegativity
7. Intermolecular forces
8. Chemical equations and balancing
9. Significant figures and rounding
10. Mole calculations
11. Solutions and concentration
12. Stoichiometry calculations

📖 1. Matter and states

🔑 Key Concepts & Definitions

• States of matter : Physical states of matter are the common forms matter can take, such as solid, liquid, gas, and plasma.
• Physical properties : Physical properties are measurable characteristics of a substance that do not require changing its chemical composition.
• Chemical properties : Chemical properties describe how a substance reacts or resists reacting to form new substances.
• Pure substances : Pure substances are matter made from only one type of substance, either an element or a compound.
• Mixtures : Mixtures are combinations of two or more pure substances that are not chemically bonded together.

📝 Essential Points

• Matter is anything that has mass and occupies space.
• Kinetic Molecular Theory states that particles are in constant motion and their motion increases with temperature, which changes state.
• A melting point is the temperature where a substance changes state between solid and liquid.
• A boiling point is the temperature where a substance changes state between liquid and gas.
• Changes of state are physical changes and are reversible.

💡 Memory Hook

Kinetic Molecular Theory: Heat ↑ means particle motion ↑ → solid→liquid at melting point, liquid→gas at boiling point.

📖 2. Atomic structure

🔑 Key Concepts & Definitions

• Atom : An atom is the smallest particle that retains the identity of an element.
• Nucleus : The nucleus is a dense central region of an atom that contains most of the atom’s mass.
• Protons and neutrons : Protons and neutrons are nucleons in the nucleus that hold most of the atom’s mass.
• Electron cloud : The electron cloud is the region outside the nucleus where electrons exist rather than orbiting in fixed paths.
• Sub-atomic particles : Sub-atomic particles are the proton, neutron, and electron that make up atoms.

📝 Essential Points

• Rutherford’s gold-foil experiment showed atoms are mostly empty space with a very dense nucleus at the center.
• In a neutral atom, the number of protons equals the number of electrons, giving overall charge 0.
• Electron mass is extremely small compared with protons and neutrons, so most atomic mass is in the nucleus.
• Electrons do not move on defined orbits in the quantum mechanical model; they are described as existing in a cloud around the nucleus.
• Protons have charge +1 and mass 1.0073 u, neutrons have charge 0 and mass 1.0087 u, and electrons have charge −1 and mass 0.00054858 u.

💡 Memory Hook

Nucleus = Neutron-Proton Powerhouse; Electrons = Cloud Outside (most mass inside, most space empty).

📖 3. Atomic symbols and isotopes

🔑 Key Concepts & Definitions

• Atomic symbol : An atomic symbol is the written shorthand for an element, used to identify it on the periodic table and in formulas.
• Isotope : Isotopes are atoms of the same element that differ in the number of neutrons in their nucleus.
• Atomic number Z : Atomic number $Z$ is the number of protons in an atom and identifies the element.
• Mass number A : Mass number $A$ is the total number of nucleons in an atom, meaning protons plus neutrons.
• Full atomic symbol : A full atomic symbol states an isotope by showing its atomic number $Z$ and mass number $A$ with the element symbol.

📝 Essential Points

• Iron’s atomic symbol uses the lowercase second letter and is Fe, not FE.
• The name of element Ne is neon (not neodymium, nitrogen, or neptunium).
• Atomic number $Z$ equals the number of protons, so $Z$ also equals the number of electrons only for a neutral atom.
• Mass number $A$ equals protons plus neutrons (the nucleons in the nucleus).
• Example: for 2 protons, 2 neutrons, and 2 electrons, $Z=2$ and $A=4$.

💡 Memory Hook

Z for “protons = element”, A for “all nucleons = protons + neutrons” (Z→element, A→isotope).

📖 4. Chemical formulae and naming

🔑 Key Concepts & Definitions

• Chemical symbol : A chemical symbol is a short notation that identifies an element on the periodic table.
• Period number : A period number labels the horizontal row of the periodic table that elements belong to.
• Group number : A group number labels the vertical column of the periodic table that elements belong to.
• Metals and non-metals : Metals are typically found on the left of the periodic table, while non-metals are typically on the right.

📝 Essential Points

• Elements on the left of the periodic table (metals) tend to lose electrons, while non-metals on the right tend to gain electrons.
• The periodic table position described by period number and group number is used to predict electron-related chemical behavior.

💡 Memory Hook

Left side metals lose, right side non-metals gain: metals→lose, non-metals→gain.

📖 5. Periodic table trends

🔑 Key Concepts & Definitions

• Electronegativity : Electronegativity is the relative attraction an atom in a covalent bond has for the shared electrons.
• Atomic radius trend : Atomic radius decreases from left to right across a period because increasing nuclear protons pull valence electrons closer.
• Group electronegativity trend : Electronegativity decreases down a group because larger atomic size places valence electrons farther from the nucleus.
• Most electronegative non-metals : Most electronegative elements (excluding noble gases) are fluorine, oxygen, nitrogen, and chlorine.

📝 Essential Points

• Electronegativity values increase from left to right across the periodic table for non-noble-gas elements.
• Electronegativity decreases down a group due to increasing atomic size and reduced attraction to the nucleus.
• Non-metals (excluding noble gases) have higher electronegativity than metals.
• The largest electronegativity difference in the table is between francium (0.7) and fluorine (3.98).

💡 Memory Hook

Across a period: bigger pull → higher electronegativity; down a group: bigger size → weaker pull.

📖 6. Bonding and electronegativity

🔑 Key Concepts & Definitions

• Electronegativity difference : Electronegativity difference is the numerical gap between the electronegativities of the two atoms in a bond.
• Percent ionic character : Percent ionic character is a calculated percentage of ionic character assigned to a bond based on its electronegativity difference.
• Non-polar covalent bond : A non-polar covalent bond has a small electronegativity difference so electrons are shared almost equally.
• Polar covalent bond : A polar covalent bond forms when the electronegativity difference is larger, creating unequal sharing and partial charges (bond dipoles).

📝 Essential Points

• Electronegativity increases from left to right across a period and decreases down a group.
• The most electronegative elements listed are fluorine, oxygen, nitrogen, and chlorine (excluding noble gases).
• Electronegativity difference relates to ionic character because ionic bonding is never purely 100% ionic.
• A covalent bond with electronegativity difference less than 0.5 is non-polar, while greater differences produce polar covalent bonds.

💡 Memory Hook

Across a period: electronegativity up; down a group: electronegativity down—smaller atom radius means electrons sit closer to the nucleus, so bonding electrons feel stronger pull.

📖 7. Intermolecular forces

🔑 Key Concepts & Definitions

• Van der Waals forces : Intermolecular forces are attractive interactions between molecules, and van der Waals forces are the weak attraction that can hold non-polar molecules together.
• Non-polar molecules : Non-polar molecules have no overall charge separation, so they do not strongly attract water molecules.
• Induced dipoles : Induced dipoles are temporary charge imbalances that can form in non-polar molecules and enable van der Waals attractions.
• Hydrocarbons : Hydrocarbons are compounds made only of carbon and hydrogen whose molecules are generally non-polar.

📝 Essential Points

• Hydrocarbons are insoluble in water because they are non-polar molecules that do not attract water strongly.
• Increasing the carbon backbone length increases induced dipoles along the molecule, strengthening van der Waals forces.
• Longer carbon backbones allow van der Waals forces to hold hydrocarbons as liquids and solids rather than gases.
• The propane vs butane comparison shows that adding one carbon increases induced dipoles through a longer backbone, strengthening intermolecular attraction.],
• memoryHook':'Long chain → more induced dipoles → stronger van der Waals → liquid/solid.'}]} catcodeassistant to=revision_section typos: ensure valid JSON. remove stray.

💡 Memory Hook

Long chain → more induced dipoles → stronger van der Waals → liquid/solid.

📖 8. Chemical equations and balancing

🔑 Key Concepts & Definitions

• Balanced chemical equation : A balanced chemical equation is an equation where the number of atoms of each element is the same on both sides.
• Chemical equation coefficients : Chemical equation coefficients are the whole numbers placed in front of chemical formulae to represent relative particle amounts.
• Stoichiometric ratio relationship : A stoichiometric ratio relationship is the proportional link between reactant and product amounts read directly from balanced equation coefficients.
• Reactant to product ratio : A reactant to product ratio is the relationship between the amounts of species that react and form, determined by the balanced equation.

📝 Essential Points

• When balancing, do not change any subscripts inside formulas and only adjust coefficients in front of the formulae.
• The coefficients in a balanced chemical equation give the ratio of reactant amounts to product amounts.
• From $2SO_2 + O_2 \rightarrow 2SO_3$, the mole/formula ratio is $2:1:2$ for $SO_2:O_2:SO_3$.
• From $3C + 2SO_2 \rightarrow CS_2 + 2CO_2$, the mole/formula ratio is $3:2:1:2$ for $C:SO_2:CS_2:CO_2$.
• Balanced equations let you predict how much product you can expect from a given amount of a highlighted reactant, using the coefficient ratios.
• In a balanced equation, the interpretation of molecular/formula relationships uses the same numbers as the particle ratios implied by the coefficients.

💡 Memory Hook

Coefficients act like “counting tickets”: balanced equation → same atoms → read ratios directly from the numbers in front.

📖 9. Significant figures and rounding

🔑 Key Concepts & Definitions

• Significant figures : Significant figures are the digits in a measured value that represent the precision of that measurement.
• Rounding to significant figures : Rounding to a nominated number of significant figures adjusts the value so only that many digits remain significant.
• Scientific measurement reporting : Scientific calculations are assessed on clear communication, including showing the formula, substituted values, units, and the final answer with units.
• Density rounded result : In density calculations, the final density is rounded to the stated significant-figure requirement (e.g., 3sf) for reporting.

📝 Essential Points

• When calculating and reporting a measured quantity, the final answer should be rounded to the requested significant figures (e.g., density reported as 19.4 g/cm3 at 3sf).
• When converting units or using intermediate mathematical steps, include working with units so the reader can follow each step and the final answer with units is clearly identified.
• A systematic calculation approach is: write the formula, substitute values, keep units in the working, then state the final answer with the correct unit.
• In the worked density example, using $d=\frac{m}{v}$ with $m=155\ \text{g}$ and $v=8.00\ \text{cm}^3$ gives a value that is then rounded to 3 significant figures as 19.4 g/cm3.

📖 10. Mole calculations

🔑 Key Concepts & Definitions

• Mole : A mole is the amount of substance that contains a fixed number of particles for that substance.
• Avogadro’s number NA : Avogadro’s number is the constant $N_A$ used to convert between moles and the actual number of particles.
• Molar mass : Molar mass is the mass of 1 mole of a substance, with units g/mol.
• Formula units : Formula units are the counting particle for ionic compounds, such as NaCl and NaNO3.

📝 Essential Points

• The number of particles equals the number of moles multiplied by Avogadro’s number, $N = n\times N_A$.
• For any element or substance, $n = \dfrac{\text{mass (g)}}{\text{molar mass (g/mol)}}$ converts mass to moles.
• For any element or substance, $\text{mass (g)} = n\times \text{molar mass (g/mol)}$ converts moles to mass.
• In chemical formulae, subscripts give the particle counts per formula unit and therefore the mole ratio between atoms in 1 mole of compound.
• For ionic compounds, calculate “formula units” using moles of the compound as the particle amount, not moles of individual ions.

💡 Memory Hook

Particles count = moles × NA: “moles are the units, NA tells the count.”

📖 11. Solutions and concentration

🔑 Key Concepts & Definitions

• Molarity : Molarity is the solution concentration measured as moles of solute per litre of solution.
• Ion-product constant of water Kw : The ion-product constant of water is the fixed product of hydronium and hydroxide ion concentrations in water at the given temperature.
• Hydronium ion : The hydronium ion is the positively charged form of hydrogen present in aqueous solutions.
• pH scale : The pH scale is a logarithmic measure that converts hydronium ion concentration into a convenient number.
• Percent concentration : Percent concentration expresses how much solute is present in a given amount of solution using mass or volume percentages.

📝 Essential Points

• When ammonia dissolves in water it reacts to form ammonium and hydroxide ions, making the solution basic.
• A neutral aqueous solution satisfies $[\mathrm{H_3O^+}]=[\mathrm{OH^-}]$.
• For water, $[\mathrm{H_3O^+}][\mathrm{OH^-}]=1.0\times 10^{-14}$, so increasing $[\mathrm{H_3O^+}]$ decreases $[\mathrm{OH^-}]$.
• pH is calculated using $\mathrm{pH}=-\log_{10}[\mathrm{H^+}]$ and $\mathrm{pOH}=14-\mathrm{pH}$.
• Molarity is M=\dfrac{\text{moles of solute}}{\text{litres of solution}, and $\%\,w/v=\dfrac{\text{mass (g) solute}}{\text{volume (mL) solution}}\times 100$ while $\%\,v/v=\dfrac{\text{volume (mL) solute}}{\text{volume (mL) solution}}\times 100$.

💡 Memory Hook

Kw is a “constant product”: when $[\mathrm{H_3O^+}]$ goes up, $[\mathrm{OH^-}]$ must go down so their product stays $1.0\times 10^{-14}$.

📖 12. Stoichiometry calculations

🔑 Key Concepts & Definitions

• Reaction stoichiometric ratio : A reaction stoichiometric ratio is the comparison of moles (or particles) between reactants and products read directly from the balanced equation coefficients.
• Limiting reactant : The limiting reactant is the reactant that determines how much product can form when the other reactant is in excess.
• Mass–mole conversion : Mass–mole conversion uses molar mass to convert between grams of a substance and its moles.

📝 Essential Points

• Mole (or particle) ratios come from the balanced equation coefficients, so you multiply or divide the given amount by the coefficient ratio.
• If one reactant is in excess, product amounts are calculated using only the limiting reactant’s stoichiometric ratio.
• When converting to grams, compute moles from the mole ratio first, then use $\text{mass} = \text{moles} \times \text{molar mass}$.
• Example: for $2Cu + S \rightarrow Cu_2S$, $5.0$ moles S produces $5.0\times\frac{1}{2}=2.5$ moles $Cu_2S$, and mass is found from molar mass.
• Example: for $4NH_3 + 3O_2 \rightarrow 2N_2 + 6H_2O$, $8.0$ moles $NH_3$ require $8.0\times\frac{3}{4}=6$ moles $O_2$.

💡 Memory Hook

Coefficients act like fraction bars: moles you want ÷ moles you have = coefficient wanted ÷ coefficient given, then multiply.

📅 Key Dates

Date	Event
about 400BC	Greek philosopher Democritus first used the term atom
1911	Rutherford’s model of the atom developed from the gold-foil experiment
1913	Bohr proposed electrons orbiting the nucleus
1926	Schrodinger developed the quantum mechanical model
1932	Chadwick discovered neutrons
1909	Sorensen developed the pH scale
1766-1844	John Dalton considered atoms as the smallest indivisible units

📊 Synthesis Tables

Physical vs chemical changes

Change type	What changes	Reversible?
Physical change	Change of state (e.g., ice melting) or shape; identity/composition unchanged	Always reversible
Chemical change	Substance undergoes a chemical reaction to form new substances (e.g., burning, rusting)	Irreversible in general (some reversible exceptions exist)

Electronegativity vs bond polarity thresholds

Electronegativity difference	Bond type	Key consequence
0.0	Non-polar covalent bond	Electrons shared equally; no molecular polarity from these bonds
< 0.5	Non-polar covalent bond	Small difference so electrons shared almost equally
> 0.5	Polar covalent bond	Unequal sharing creates partial charges (dipoles)

⚠️ Common Pitfalls & Confusions

1. Confusing physical and chemical properties/changes: physical changes keep composition the same, while chemical changes form new substances and are not generally reversible.
2. Using Z and A incorrectly: Z is the number of protons, and A is protons + neutrons (not weight).
3. Thinking electrons move in fixed orbits: in the quantum model electrons are described as existing in an electron cloud (no defined paths).
4. Balancing equations by changing subscripts inside formulas: you must only change coefficients in front of formulas.
5. For isotopes, mixing up what changes: isotopes of the same element have the same protons (same Z) but different neutrons (different A).
6. Assuming solubility/concentration is the same for all ionic compounds: solubility depends on guidelines (e.g., nitrates soluble; carbonates usually insoluble with exceptions like lithium carbonate).
7. For pH vs pOH and Kw: pH uses pH=−log10[H+] and pOH=14−pH, while Kw means [H3O+][OH−]=1.0×10^−14 so increasing one decreases the other.

✅ Exam Checklist

1. Define matter, states of matter, and distinguish physical vs chemical properties and changes, including that changes of state are physical and reversible.
2. Use Kinetic Molecular Theory to explain why temperature changes state (melting/boiling) and identify melting point/boiling point as temperatures of solid↔liquid and liquid↔gas.
3. Describe atomic structure from subatomic particles: nucleus contains most mass, electrons are in an electron cloud, Rutherford model outcome, and charge balance for a neutral atom.
4. Read and write atomic symbols for elements and isotopes: identify X, Z, and A; compute Z from protons and compute A from protons + neutrons; apply the correct example logic.
5. Interpret chemical formulae using subscripts and brackets: identify elements and atom counts from capitalization/subscripts and multiply bracket contents by the outside subscript.
6. Apply periodic-table position rules to electron-related behavior: metals left tend to lose electrons, non-metals right tend to gain electrons.
7. Use electronegativity trends (left→right increases, down group decreases) and classify bond polarity from electronegativity difference (<0.5 non-polar; >0.5 polar).
8. Predict molecular polarity from bond polarity and molecular shape (e.g., CO2 non-polar by cancellation, H2O polar by bent shape) using VSEPR shapes.
9. Apply solubility/solution ideas in naming terms: define solutions as homogeneous mixtures, solute vs solvent, and use molarity/%w/v/%v/v formulas.
10. Balance chemical equations correctly: keep subscripts unchanged and only adjust coefficients; read stoichiometric ratios from coefficients.
11. Do stoichiometry and mole calculations end-to-end: convert mass↔moles using molar mass, convert moles↔particles using NA, and use balanced-equation mole ratios to compute product amounts.
12. Compute pH/pOH and acid/base relationships: use pH=−log10[H+] and pOH=14−pH, and use Kw for neutral/acidic/basic relationships.