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The concept that matter is made up of tiny discrete particles has evolved over centuries through philosophical ideas, experimental observations, and scientific reasoning. Understanding this historical development helps us appreciate how modern chemistry emerged.
In ancient times, philosophers debated whether matter was continuous or made up of indivisible particles.
Despite Aristotle's influence, the idea of atoms persisted in some circles, laying groundwork for future scientific inquiry.
In the early 19th century, John Dalton revived the atomic concept with scientific evidence, formulating the first modern atomic theory.
This theory explained several experimental laws:
Dalton’s atomic theory was supported by quantitative experiments:
Discovery of subatomic particles refined the atomic theory:
The gold foil experiment demonstrated that most alpha particles passed through the foil, but some were deflected at large angles, indicating a concentrated positive nucleus.
Today, the particulate nature of matter is understood at the atomic and subatomic levels:
This modern view explains chemical bonding, molecular structure, and the properties of matter in detail.
This timeline summarizes the key milestones from ancient atomism to modern quantum theory.
Difficulty: ★☆☆☆☆ (Easy)
Problem: According to Dalton’s atomic theory, what happens to atoms during a chemical reaction?
Solution: Dalton’s theory states that atoms are neither created nor destroyed during a chemical reaction. Instead, atoms are rearranged to form new compounds. Therefore, the number of atoms remains the same, only their combinations change.
Difficulty: ★★☆☆☆ (Moderate)
Problem: Carbon and oxygen form two compounds: CO and CO\(_2\). If 12 g of carbon combines with 16 g of oxygen to form CO, how much oxygen combines with 12 g of carbon to form CO\(_2\)? Verify the law of multiple proportions.
Solution:
Given: Mass of carbon = 12 g (constant)
For CO, oxygen = 16 g
For CO\(_2\), oxygen = ?
Molecular formula CO\(_2\) has twice the oxygen atoms as CO, so oxygen mass = \(16 \times 2 = 32\) g.
Ratio of oxygen masses combining with fixed carbon mass:
\[ \frac{32}{16} = 2 \]
This is a simple whole number ratio, confirming the law of multiple proportions.
Difficulty: ★★★☆☆ (Intermediate)
Problem: A compound contains 70% iron and 30% oxygen by mass. If the compound’s formula is Fe\(_2\)O\(_3\), calculate the relative atomic mass of iron.
Solution:
Assume 100 g of compound:
Mass of Fe = 70 g
Mass of O = 30 g
Moles of oxygen atoms:
\[ n_O = \frac{30}{16} = 1.875 \]
From formula Fe\(_2\)O\(_3\), mole ratio Fe:O = 2:3 = 0.6667
Let relative atomic mass of Fe be \(M_{Fe}\). Moles of Fe:
\[ n_{Fe} = \frac{70}{M_{Fe}} \]
Using mole ratio:
\[ \frac{n_{Fe}}{n_O} = \frac{2}{3} \Rightarrow \frac{70 / M_{Fe}}{1.875} = 0.6667 \]
Simplify:
\[ \frac{70}{M_{Fe} \times 1.875} = 0.6667 \Rightarrow M_{Fe} = \frac{70}{1.875 \times 0.6667} \approx 56 \]
The relative atomic mass of iron is approximately 56, matching the known value.
Difficulty: ★★★☆☆ (Intermediate)
Problem: Why did most alpha particles pass through the gold foil undeflected in Rutherford’s experiment, while some were deflected at large angles?
Solution: Most alpha particles passed through because atoms are mostly empty space. The few particles deflected at large angles encountered the small, dense, positively charged nucleus, which repelled the positively charged alpha particles, causing deflection.
Difficulty: ★☆☆☆☆ (Easy)
Problem: Name the three main subatomic particles and state their charges.
Solution:
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