Orbitals of the Outermost Electron

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Balancing Chemical Equations Worksheet (with answer key)

 Balance the following chemical equations by writing the correct coefficients in the blanks provided.

1. __ H₂ + __ O₂ ⟶ __ H₂O

2. __ S₈ + __ O₂ ⟶ __ SO₂

3. __ Al + __ O₂ ⟶ __ Al₂O₃

4. __ Na + __ Cl₂ ⟶ __ NaCl

5. __ CH₄ + __ O₂ ⟶ __ CO₂ + __ H₂O

6. __ Fe + __ H₂O ⟶ __ Fe₃O₄ + __ H₂

7. __ P₄ + __ O₂ ⟶ __ P₂O₅

8. __ C₃H₈ + __ O₂ ⟶ __ CO₂ + __ H₂O

9. __ K + __ H₂O ⟶ __ KOH + __ H₂

10. __ H₂O₂ ⟶ __ H₂O + __ O₂

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Answer Key

1. 2 H₂ + 1 O₂ ⟶ 2 H₂O

2. 1 S₈ + 8 O₂ ⟶ 8 SO₂

3. 4 Al + 3 O₂ ⟶ 2 Al₂O₃

4. 2 Na + 1 Cl₂ ⟶ 2 NaCl

5. 1 CH₄ + 2 O₂ ⟶ 1 CO₂ + 2 H₂O

6. 3 Fe + 4 H₂O ⟶ 1 Fe₃O₄ + 4 H₂

7. 1 P₄ + 5 O₂ ⟶ 2 P₂O₅

8. 1 C₃H₈ + 5 O₂ ⟶ 3 CO₂ + 4 H₂O

9. 2 K + 2 H₂O ⟶ 2 KOH + 1 H₂

10. 2 H₂O₂ ⟶ 2 H₂O + 1 O₂

Lewis Structures Practice

Instructions

Draw the Lewis structure for each of the molecules listed.

1. H₂S 
2. CF₄ 
3. NH₃ 
4. CS₂ 
5. Br₂ 
6. CBr₄ 
7. HCN 
8. SO₂ 
9. PBr₃ 

Lewis Structures

Lewis structures are visual representations of the valence electron arrangement within molecules and polyatomic ions. They help us understand bonding and predict molecular geometry. The fundamental goal is to arrange atoms and valence electrons so that each atom achieves a stable electron configuration, typically an octet (eight valence electrons), resembling a noble gas. The first step in constructing a Lewis structure is to count the total number of valence electrons for all atoms in the molecule or ion. For ions, remember to add electrons for negative charges and subtract for positive charges.

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Next, you need to determine the central atom. This is usually the least electronegative atom (excluding hydrogen, which is always terminal). If there are multiple atoms of the same element, the single atom of that element is often central (e.g., carbon in CH₄). Once the central atom is chosen, draw a skeletal structure by connecting all other atoms (terminal atoms) to the central atom with a single bond. Each single bond represents two shared valence electrons. Subtract these bonding electrons from your total valence electron count.

Now, distribute the remaining valence electrons as lone pairs to satisfy the octet rule for all terminal atoms first. Hydrogen atoms only need two electrons (a duet) and are satisfied with their single bond. After satisfying the terminal atoms, place any leftover electrons as lone pairs on the central atom.

Finally, check if all atoms have an octet (or a duet for hydrogen). If the central atom (or any other atom) does not have a complete octet, you may need to form multiple bonds. Move one or more lone pairs from a terminal atom to form double or triple bonds with the central atom until all atoms achieve an octet. Remember to recalculate formal charges if you are evaluating different resonance structures, though for basic construction, ensuring octets is the primary goal.

Periodic Table Trends (Electronegativity, Ionization energy, and Atomic radius)

The periodic table arranges elements to reveal predictable patterns in their properties, known as periodic trends. These trends are primarily governed by two factors: the number of electron shells (energy levels) and the effective nuclear charge, which is the pull that the outermost electrons feel from the nucleus. Understanding these core principles allows us to predict an element's behavior based on its position on the table. Three of the most important trends are atomic radius, ionization energy, and electronegativity.

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Atomic Radius 

Atomic radius is a measure of the size of an atom. As you move down a group on the periodic table, the atomic radius increases. This is because each step down adds a new electron shell, placing the outermost electrons further from the nucleus. Moving from left to right across a period, the atomic radius decreases. Although electrons are being added, they are filling the same energy level, while protons are also being added to the nucleus. This stronger nuclear charge pulls the electron cloud in more tightly, shrinking the atom.

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Ionization Energy 

Ionization energy is the energy required to remove an electron from an atom. This trend is inversely related to atomic radius. As you move down a group, the ionization energy decreases. The outermost electrons are farther from the nucleus and are shielded by inner electrons, making them easier to remove. Conversely, moving across a period, ionization energy increases. The greater effective nuclear charge holds the electrons more tightly, so it takes significantly more energy to strip one away.

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Electronegativity 

Electronegativity describes an atom's ability to attract shared electrons in a chemical bond. Moving down a group, electronegativity decreases. With more electron shells, the nucleus is farther from the shared bonding electrons and has a weaker pull on them. As you move across a period, electronegativity increases due to the rising nuclear charge, which more powerfully attracts the bonding electrons. Fluorine is the most electronegative element, while noble gases are typically assigned an electronegativity of zero as they rarely form bonds.

You can also find this explained in this video. 

Historical Models of the Atom

Our understanding of the atom has evolved significantly over time, starting from a simple, indivisible sphere to the complex quantum model we use today. Early ideas began with ancient Greek philosophers like Democritus, who first proposed that matter was composed of tiny, indestructible particles called "atomos." However, it was John Dalton in the early 1800s who established the first scientific atomic theory. Dalton's model envisioned the atom as a solid, indivisible billiard ball, suggesting that all atoms of a particular element were identical and that chemical reactions were simply the rearrangement of these atoms.

This "billiard ball" model was revolutionized in 1897 by J.J. Thomson's discovery of the electron. His experiments with cathode ray tubes revealed that atoms contained negatively charged particles. This led to the "plum pudding" model, where the atom was depicted as a positively charged sphere with negatively charged electrons embedded within it, much like plums in a pudding. This was a major step forward, as it introduced the concept that atoms were, in fact, divisible and had an internal structure.

The plum pudding model was soon replaced by Ernest Rutherford's nuclear model in 1911. His famous gold foil experiment showed that the atom is mostly empty space with a tiny, dense, positively charged nucleus at its center, with electrons orbiting it. A few years later, Niels Bohr refined this model by proposing that electrons travel in specific, fixed orbits or "energy levels" around the nucleus, and they could jump between these levels by absorbing or emitting energy. While the Bohr model has since been superseded by the modern quantum mechanical model, which describes electrons in terms of probability clouds rather than fixed orbits, these earlier models were crucial stepping stones that progressively built our accurate and detailed picture of the atom.