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2 years ago

What are examples of models

1 answer:

6 0

Here are 3 main types of models are ( Physical, Mathematical, and Conceptual. Models have limitations but are useful and can be changed based on a new evidence.

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The reaction ag+(aq) + cl−(aq) ↔ agcl(s) has an equilibrium of 1020 (keq = 1020). if you have a beaker containing 1 liter of wat

lord [1]

When Q is equal the initial concentration of the products / the initial concentration of the reactants.

so, Q = [Ag]*[Cl-] and we neglected [AgCl] as it is solid

∴ Q = 10^-6 * 10^-5

= 10^-11

now we will compare the value of Q with the value of Keq:

when Q = Keq so, the system is in equilibrium

and when Q > Keq so, the reaction will go forward (shift to right) to achieve equilibrium.

and when Q< Keq so, the reaction will go left (shift to left) to achieve equilibrium.

when Q = 10^-11 and Keq = 10^20

∴Q< Keq

and the reaction will shift to left.

3 0

2 years ago

Suppose 1.87g of nickel(II) bromide is dissolved in 200.mL of a 52.0mM aqueous solution of potassium carbonate.

Reika [66]

Answer:

0.0428 M

Explanation:

Because we're asked to calculate the molarity of nickel(II) cation, we need to determine all sources for that species, in this case, all Ni⁺² comes from the nickel(II) bromide solid (NiBr₂).

We use the molecular weight of NiBr₂ to calculate the moles of Ni:

1.87 g NiBr₂ ÷ 218.49g/mol * (1molNi⁺²/1molNiBr₂) = 8.55x10⁻³ mol Ni⁺²

Then we divide the moles by the volume in order to calculate the concentration:

8.55x10⁻³ mol Ni⁺² / 0.200 L = 0.0428 M

7 0

2 years ago

A compound has been shown to contain 14.5 percent hydrogen and 85.5 percent carbon by mass.

Sedaia [141]

Answer:

Option d. Its empirical formula is CH₂.

Explanation:

The percent composition of the compound allow you to calculate the empirical formula of the compound but is not enough to calculate either the molar mass or the molecular formula. So, since now you can discard options b. and c.

Telling that it is a hydrocarbon (option e.) is true but very vague compared with finding the empirical formula. So, you can also discard the option e.

The fact that the product has a triple bond cannot be concluded from the percent composition, you should find the molecular formula to assert whether it contains or not a triple bond. So, you could discard option a., which lets you only with choice d.

Let us find the empirical formula to be certain that it is CH₂.

1. First, assume a basis of 100 g of compound:

H: 14.5% × 100 g = 14.0 g

C: 85.5% × 100 g = 85.5 g

2. Divide each element by its atomic mass to find number of moles:

H: 14.0 g / 1.008 g/mol = 14.38 mol

C: 85.5 g / 12.011 g/mol = 7.12 mol

3. Divide both amounts by the smallest number, to find the mole ratio:

H: 14.38 mol / 7.12 mol ≈ 2
C: 7.12 mol / 7.12 mol = 1.

Hence, the ratio is 2:1 and the empirical formula is CH₂.

7 0

2 years ago

Pleaseeee helpppp thank youu

Gemiola [76]

Answer:

<h3>Answer: well </h3>

b. a type of gas is evolved ( hydrogen gas )

5 0

2 years ago

Consider the following reaction:C2H4(g) + F2(g) -----------> C2H4F2(g) Delta H = -549 kJEstimate the carbon-fluorine bond ene

frutty [35]

Answer:

Bond energy of carbon-fluorine bond is 485 kJ/mol

Explanation:

Enthalpy change for a reaction, is given as:

$\Delta H_{rxn}=\sum [n_{i}\times (E_{bond})_{i}]-\sum [n_{j}\times (E_{bond})_{j}]$

Where $(E_{bond})_{i}$ and $(E_{bond})_{j}$ represents average bond energy in breaking "i" th bond and forming "j" th bond respectively. $n_{i}$ and $n_{j}$ are number of moles of bond break and form respectively.

In this reaction, one mol of C=C, four moles of C-H and one mol of F-F bonds are broken. One mol of C-C bond, four moles of C-H bonds and two moles of C-F bonds are formed

So, $-549kJ=(1mol\times 614kJ/mo)+(4mol\times E_{C-H})+(1mol\times 154kJ/mol)-(1mol\times 347kJ/mol)-(4mol\times E_{C-H})-(2mol\times E_{C-F})$

or, $-549kJ=(1mol\times 614kJ/mo)+(1mol\times 154kJ/mol)-(1mol\times 347kJ/mol)-(2mol\times E_{C-F})$

or, $E_{C-F}=485kJ/mol$

So bond energy of carbon-fluorine bond is 485 kJ/mol

8 0

2 years ago